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Chairman,  Drpmtmrnl  of  j'hyuftil  Science 
Jnmr\  Wutli\nn  ////,'/<  .SV hoitl,  AV/i  York  (Iity 






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[\K\\  \\om.l)  OF  r.llhMISTRYua*  In-r  pulili^linl  in  lU3i 
and  ri»ni|ilrtrh  ir\i->rd  in  L(MJ  and  11U7.  Minor  icvi^ion- 
lia\«-  IMTII  madr  r\n\  t\\o  or  llirrc  \«MI-  to  krr|»  (lu*  hook 
up-(o-dalr  'l'ln-«  I1).").")  nlition  i>«  .in  fxtcn^ue  rr\iNioii  'I'ht-  tr\t 
IKKS  IM>CII  Kirp-l\  rruiitlcn,  nr\vl\  lio^i^ncd.  nculx  illn-lratrd. 

Format  Design  .  .  .  OliMe  Wlntnrv 
Srratt  hhourd  Drau  ings  .  .  .  Rrno  Mai  tin 
NT  in  Ba^kcrnllc 
Printed  by   lMiin|iton  l'rr*s 



.  For  success  and  satisf  action  in  modern  living,  tin*  citi/en 
must  understand  and  appreciate  the  scientific  aspects  of  his 
environment  and  the  role  ol  science  in  the  development  of 
civili/ation.  lie  must  hase  his  everydav  thinking  and  action 
on  the  best  information  available  with  a  full  reali/ation  of 
why  the  methods  of  science  are  superior  to  other  methods  of 
obtaining  information.  Finally,  the  citi/en  must  be  ready  to 
use  scientific  knowledge  lor  the  <M)od  of  all.  M  \v  \\'ORI,|)  or 
CIII.MISTKY  is  designed  to  help  students  achieve  these  objec- 
tives. It  is  a  textbook  for  yoimjj;  people  who  are  learning  to  be 

This  1  ()r).")  edition  of  MAV  \VOKI.D  01-  CIIMMISTRY  lias  been 
brought  abreast  of  recent  chemical  developments  and  dis- 
coveries. Striking  advances  have  been  made  in  nuclear  energy, 
petrochemistry,  metallurgy,  textiles,  and  plastics.  The  chapters 
dealing  with  these  subjects  have  been  tlionni^lHy  revised  to 
take  these  changes  into  account. 

The  basic  plan  ol  the  book,  which  has  been  so  widely 
accepted,  lias  not  been  altered.  Hut  within  that  framework 
many  changes  have  been  made.  Diagrams  have  been  com- 
pletely redrawn  and  enlarged.  New  illustrations  have  been 
provided.  The  index  lias  been  expanded  to  make  the  text  an 
even  more  effective  source  book.  Despite  the  addition  of  new 
materials,  however,  the  hook  as  a  whole  lias  been  shortened. 
Perhaps  most  important  of  all.  the  chapters  dealing  with 
the  basic  theory  and  mathematics  of  chemistry  have  been 
rewritten.  Formulas,  valence,  equations,  and  problems  have 
been  £iven  a  new  and  expanded  treatment.  The  chapters  pre- 
senting the  electron  theory,  ioni/ation.  and  the  study  of  teases 
have  been  rewritten  to  achieve  a  maximum  of  clarity.  While 
the  concepts  of  chemistry  themselves  cannot  be  simplified. 


thry  can  he  made  r.isiri  to  understand  h\  hcttcT  arrangrmrnt 
and  tfn.ilci  clanlx  of  laniiua^c'  Glarit)  of  wilting  has  been 
.L  m.ijoi  140.1!  of  this  substantial  levisinn  of  NKW  WORLD  OF 


'I  he  boc»k  has  been  \\iitten  \\ith  the1  puipose  of  getting  idea* 
acrms  to  the  \tudcnL  With  the  exception  of  ncccssaix  teehnical 
trims,  tin  \o(,ibulai\  consists  of  familial  words.  The  style  is 
duett,  tin  sentences  and  paia<;iaphs  are  short.  Much  of  the 
illiistiatm  mateiial  has  been  draxxn  fiom  the  student's  own 
enviionmc  nt  (lonsumii  aspects  of  che-mistix  and  the  them- 
isti\  of  common  things  aie  stiessed  at  the  same  time  that 
ic(|iiiienie.its  of  model n  couises  of  stud\  and  examining  boards 
aie  iulh  met. 

Select!  d  matenals  from  the  lnstor\  of  c  hemistr\  arouse 
inteiest  .ind  enable  the  student  to  sec  how  fundamental  scien- 
tific ideas  ha\e  developed  and  expanded.  Thc\  revc\il  that 
rliemic  .il  thioiies  ,md  principles  arc*  descriptive  generali/ations 
of  ni. ill's  oinv  ing  i  xpiMieni  i  in  tr\mi>  to  undei stand  his  world. 
FinalK,  the1  hi^tonc  materials  hel])  studrnts  scu-  hem  scientific 
discoxm  has  aflerte-d  modem  civili/ation  and  hem  the1  needs 
of  sot  iet\  c'onstantK  stimulate*  s(  lentific  rest-. in  h 

The*  authoi  \\islics  to  expiess  his  thanks  to  the  mail} 
teachcTs.  diemisK  and  industrial,  e'diu  and  govc-rn- 
menttil  or^.un/ations  that  have1  generousU  made  suggc-stions, 
chicked  the  acciuacA  of  mateiials.  and  provided  illustiations 
and  data  foi  this  nc\\  edition.  He  aNo  \\ishes  to  ac'kncm ledge* 
his  indebtedness  to  his  son.  I)i  Lionel  F  Jaffe-,  and  to  Lrc 
I)ei»hton  and  John  H  \Villi.iniscm  of  Silver  Buidett  (Company 
\\liosi-  i-aieliil  plaiinui!*.  sound  judgmc^nl,  and  nrvcr-f ailing 
rntliiisia^ni  h.ixe  made  this  c-oopeiative-  efTort  e-xtie-me^h  pleas- 
ant and,  it  is  hoped,  eflrctive 

Ft  nun  id  Jaffe 


1.  MATTER  and  Its  Changes  I 

What    the    \\oild    is    made    of    .uul    ho\\     M  icntilu     methods    expand 
our  undci^tandmg. 

2.  OXYGEN:  Earth's  Most  Abundant  Element  24 

Distovei),    piopeities,    and    uses   of   the   gas   \\lmli    is    the    bicath    of 

3.  HYDROGEN:  Lightest  of  the  Elements  44 

The   piimat\    stnlT  of  \\hkh  tlie  entne  uni\eise  is  l)inlt. 

4.  WATER:  Most  Common  Liquid  Compound  60 

The  thiillmg  discoveiy  of  its  composition,   its  ptopeitics,  and    uses. 

5.  ATOMS:  Bricks  of  the  Universe  75 

\  Quaker  sc  hoolteat  her  explains  t\\o   laus  of  <hemisti\    b\    means 
of  his  atoms 

6.  FORMULAS:  The  Chemist's  Abbreviations  85 

Rules  on  the  untin^  of  chemical   foimulas. 

7.  ATMOSPHERE:  The  Ocean  of  Air  97 

I  low  the  components  of  the  air  \\ere  discovered  and   utili/ed. 

8/  EQUATIONS:  Shorthand  of  Chemistry  114 

Rules  and   aids   in   the   halancing  of  chemical    equations 

9.    MATHEMATICS  of  Chemistry  124 

Meaning  of  atomic  wriijhts  and  how  some   pioblems 
aie  solved. 

10.   CHLORINE  and  the  Halogen  Family  137 

A    closelv    related   group   of   elements — fluorine,    c  hloi  me,    bromine, 
and  iodine. 

I  L    ELECTRONS  and  Other  Particles  154 

Origins,  development,  and  chemical  usefulness  of  the  latest  theory 
of  the  structure  of  the  atom. 


12.  NUCLEAR  ENERGY  at  Last!  176 

M.m  hn.ilU  p<netiates  the  <  enter  ot  the  atom  and  produces  nu<  lear 

13.  ACIDS:  Hydrochloric  Acid,  a  Typical  Acid  195 

PiopeitKs   .iiid   g<nei,d    method   of    piep.ii  m«j   a<  ids    UK  hiding   IIF 

14.  BASES:  Sodium  Hydroxide,  a  Typical  Base  206 

Nrutr  ah/at  ion   of    bases   b\    .Kids,   and   salt    formation 

15.  SOLUTIONS:  Water,  the  Universal  Solvent  217 

Kinds  of  solutions,  distilled  and  hid:\    \\atei 

16.  IONS  and  Dissociation  231 

\\n\\   disputation  explains  stieimth  ol   acids,  ludiol\Ms,  electrolysis, 

17.  AMMONIA  and  Reversible  Reactions  250 

I  lie  lommeiu.d   s\nthesis  of   aninioma      a   milestone  in   <hemistr\ 

18.  NITRIC  ACID  and  Nitrogen  Compounds  262 

1  he    (omiiKKial    s\nthesis    of    nitiit     .Kid,    and    mttogeii    fixation 

19.  MOLECULES:  Avogadro's  Hypothesis  276 

Imliidmu     applications     lo     the     solution     of     \\ni;lit-\oltinir     and 
sliai^hl-\  olume    pioblenis 

20.  SULFUR  and  Hydrogen  Sulfide  289 

I  hen    piodmtion,   piopeities,  and  uses,   .IK  hiding  the  manuf.ieture 
of   mat(  lies 

21.  SULFURIC  ACID:  The  Fundamental  Acid  304 

Its  (   piepaiation,   piopeities,  and  mam    uses. 

22.  ALLOTROPIC  CARBON:  Key  Source  of  Energy  320 

Piopeities  and  uses  \\ith  spei  attention  to  coal 

23.  CARBON  DIOXIDE:  Gas  of  Life  and  Decay  339 

Including  caibonic  aeid,  its  salts,  and  baking  poudeis. 

24.  CARBON  MONOXIDE  and  Other  Gaseous  Fuels          352 

Including  water  gas,  producer  gas.  coal  gas,  natuial  gas,  acctvlenr 

25.  METALS  and  Their  Chemical  Activity  369 

Including  the   piepaiation,   piopcities,   uses,   and   detettion   In    the 
spectroseope  of  sodium,  potassium,  .ind  lithium. 


26.    ALUMINUM:  Most  Common  ol  Light  Metals  387 

I  IK  llldllli!    MIMIC    <>t     Its    compounds,    .Hid    j|s0    {\\c     Ill.t 

27.  IRON  and  Steel  404 

Intituling  .1  ulmipse  of  the  ne\\   linn/mis  in  the  steel  mdiis(i\. 

28.  COPPER:  Nerves  of  the  Machine  Age  423 

Its  met.  il  hit  <j,\    and  uses,  iiu  hiding  coppei    siill.itr. 

29.  OTHER  METALS  and  Their  Uses  436 

Including  the  nxe  of  the  M>-<.ilh(i  iaie  metals. 

30.  FERTILIZERS  and  Salts  of  Sodium  464 

IIK  lulling  the  solution  ol  oin   pot.ish  piohlcin. 

31.  CALCIUM:  Its  Common  Compounds  480 

u  supc  iphiisph  itcs  .UK!  \\.itrr 

32.  IRON:  Some  Special  Compounds  496 

I  IK  ludinu  then   uses  in  inks  .ind  l)luc|)iint  ]).»per. 

33.  GLASS  and  Some  Silicon  Compounds  505 

IIH  liiciinu  ,dso  some  hoTon  (  on  i  pounds  MM  h  .is  borax. 

34.  HYDROCARBONS  and  Their  Derivatives  522 

n^  the  s|)C(  t.n  nl.ii    iis<-  of    petioi  heinistix      .1   new 

35.  ALCOHOL  and  Other  Organic  Compounds  547 

Anothei  i*limp\r  of  th<-  \\cnlcl  of  the  OII^.IIIM   <  hnnist. 

36.  FOODS  and  Chemotherapy  566 

Chemistry  .it  \\oik  in  the  ser\ue  of  m,m\  he.dth 

37.  FIBERS  AND  PLASTICS:  Textiles,  Paper,  and  Dyeing      591 

Including  the  ne\\  \\oild  of  in.iii-ni.idt 

38.  COLLOIDS:  The  Colloidal  State  of  Matter  610 

The  rhcnustiy  of  tiny  paitic  Irs     drops.  fiLiincnts,  L^i.uns,  and  films 

39.  LIGHT:  Its  Chemical  Effects  623 

Including  the  principles  of  sunple  photnt^i.tphy 


Simplest  formula  and  true  formula  di  trniimmc;  atomic  s\eit^ht  ex- 
perimentally, temperature  s<ale  <on\ersjon,  use  of  Hoyle's  Jaw  and 
(ihaile^  la\>. 



.  .  .  We  must  trust  in  nothing  but 
facts.  These  are  presented  to  us  by 
nature  and  cannot  deceive.  We  ought 
in  every  instance  to  submit  our 
reasoning  to  the  test  of  experiment. 
It  is  especially  necessary  to  guard 
against  the  extravagances  of  imagina- 
tion which  incline  to  step  beyond 
the  bounds  of  truth.  Antoine  Lau- 
rent Lavoisier,  1743-1794 

Progress  through  scientific  knowledge.  A  play,  presented  in  New 
York  City,  is  seen  and  heard  instantly  and  simultaneously  in  millions 
of  homes  throughout  the  nation,  as  clearly  as  if  its  vast  audience 
were  actually  seated  in  the  studio.  An  airplane  takes  off  from  an 
airfield  in  California  and  within  a  few  minutes  is  screaming  through 
the  stratosphere  60  thousand  feet  above  the  earth  at  more  than  700 
miles  per  hour.  In  an  Alabama  farmhouse,  a  country  doctor  diag- 
noses his  patient's  illness  as  pneumonia,  and  assures  the  family  that 
it  is  quickly  curable  by  the  administration  of  certain  drugs.  In  an 
isolated  group  of  islands  in  the  South  Pacific,  a  bomb  explodes  with 
a  force  equal  to  hundreds  of  thousands  of  tons  of  dynamite.  In  Iowa, 
a  housewife  cooks  a  salt-water  fish  caught  by  her  husband  during  the 
family  vacation  six  months  before  and  two  thousand  miles  away. 

A  generation  ago,  any  of  these  occurrences  would  have  been  con- 
sidered miraculous.  Yet  today  they  arouse  no  unusual  public  excite- 
ment, for  they  are  no  longer  extraordinary  events.  To  us  they  are 
familiar  happenings,  representative  of  what  we  call  the  progress  of 

Progress  has  been  taking  place  in  some  measure  ever  since  man 
appeared  on  earth.  But  within  the  past  two  or  three  centuries, 



civilization  has  progressed  much  more  rapidly  than  in  all  the  previ- 
ous thousands  of  years  of  man's  history,  with  the  most  radical  changes 
occurring  within  the  past  few  decades.  In  the  years  ahead,  civilization 
will  continue  to  progress.  More  of  the  "incurable"  diseases  will  be 
conquered;  methods  of  transportation  and  communication  will  be 
further  improved;  the  force  which  gives  the  atomic  and  hydrogen 
bombs  their  great  destructive  power  will  be  harnessed  for  man's 
welfare;  our  everyday  lives  will  be  made  safer  and  more  com- 

What  enables  us  to  perform  acts  today  which  were  unheard  of 
just  a  few  short  years  ago?  How  may  we  be  certain  that  the  future 
will  bring  forth  new  wonders?  The  answer  to  both  of  these  questions 
lies  in  the  knowledge  we  now  possess  concerning  the  nature  and 
behavior  of  those  things  which  make  up  the  entire  world  and  all  its 
living  creatures.  In  short,  the  answer  lies  in  our  scientific  knowledge. 

Scientific  knowledge  is  not  something  which  has  been  created 
recently.  It  has  been  gathered  by  many  men,  known  and  unknown, 
during  all  of  man's  centuries  on  earth.  Just  as  a  snowball  rolling 
downhill  is  at  first  small  and  slow  moving,  but  in  time  increases  both 
its  si/e  and  speed,  so  it  has  been  with  scientific  knowledge.  In  his 
early  days  on  earth,  man  knew  a  limited  number  of  scientific  facts, 
increasing  the  number  with  painful  slowness  as  the  centuries  went  by. 

But  gradually,  almost  imperceptibly,  the  number  of  facts  became 
larger  and  their  rate  of  discovery  quickened.  Within  the  past  few 
hundred  years,  particularly  during  the  twentieth  century,  they  have 
been  acquired  at  a  prodigious  rate.  New  scientific  advances  are  an- 
nounced almost  daily.  A  rolling  snowball  must  finally  come  to  rest 
at  the  foot  of  a  hill,  but  there  is  no  indication  that  scientific  knowl- 
edge will  ever  cease  to  grow.  Nor  does  there  appear  to  be  any  limit 
to  the  progress  civilization  can  make  through  the  application  of  this 

Linde  Air  Products  Company 

(left)  The  highly-trained 
professional  chemist  makes 
many  important  contribu- 
tions to  civilization. 

(right)  The  lessons  of  the 
student  chemical  labora- 
tory find  many  useful  ap- 
plications in  daily  life. 


Chemistry  has  helped  make  a  new  world.  No  single  science  is 
entirely  responsible  for  our  modern  civilization.  But  chemistry  ranks 
high  among  those  which  are  considered  the  most  important  because 
it  is  a  basic  science^  essential  to  virtually  every  scientific  study  regard- 
less of  its  nature.  In  the  manufacture  of  every  product  of  our  great 
industrial  civilization,  chemistry  plays  ajvital  role. 

The  work  of  the  chemist  affects  all  of  us  —  the  physician,  the 
dentist,  the  engineer,  the  public  health  employee,  the  soldier,  the 
gas  station  attendant,  the  clerk,  the  housewife,  the  worker  of  factory, 
jninc,  or  farm,  the  schoolboy  and  the  schoolgirl. 

While  the  achievements  of  modern  chemists  seem  at  times  to  be 
the  result  of  some  kind  of  magic,  nothing  could  be  further  from  the 
truth.  Chemical  ^magic'^is  the  result  of  years  of  study,  hard  work. 
and  struggle;  of  burning  the  midnight  oil  when  it  seems  a  problem 
has  no  solution;  and,  finally,  of  a  thrilling  moment  when  everything 
hinges  on  one  more  experiment  and  the  world  spins  'round  in  a  test 
tube  or  flask.  Nor  are  all  experiments  rewarded  with  striking  suc- 
cess. Sometimes  they  fail  and  the  chemist  must  begin,  wearily,  but 
with  determination,  to  retrace  his  weeks  or  months  ol  work  in  an 
attempt  to  discover  where  he  went  astray. 

Chemistry  is  an  intensely  interesting  and  rapidly  changing  field 
and,  like  all  sciences,  its  roots  extend  far  back  into  history.  Today's 
chemist  owes  a  great  debt  to  the  men  who,  over  many  centuries, 
added  to  the  knowledge  he  now  uses  in  performing  his  "miracles." 
The  lives  of  those  men,  how  they  accumulated  information  and  how 
that  information  is  applied  in  modern  chemistry  is  a  thrilling  story, 
filled  with  suspense  and  drama.  It  is  the  story  told  in  NKW  WORLD  01- 


Although  we  live  in  the  "scientific  age,"  we  still  see,  everywhere 
in  the  world,  superstitions  and  prejudices,  hunger  and  disease.  To 
correct  these  conditions,  we  must  use  in  the  outside  world  what  we 
learn  in  the  classroom.  We  must  base  our  actions  upon  scientific 
knowledge,  always  bearing  in  mind,  however,  that  we  know  little 
with  absolute  finality. 

Board  of  Education,  City  of  New   York 


We  must  also  remember  that  science  is  a  two-edged  sword  which 
may  be  used  either  to  serve  man  or  to  destroy  both  him  and  his 
^ivbrks.  Although  the  spirit  of  science  is  essentially  democratic  and 
constructive,  we  alone  can  prevent  its  becoming  an  oppressively 
tyrannical  and  horribly  destructive  weapon.  Not  many  of  us  may 
ever  become  professional  scientists,  the  men  and  women  who  work 
in  the  great  laboratories  of  industry,  government,  and  education. 
But  we  can  all  become  scientists  in  an  even  broader  sense. 

We  can  act  on  the  most  dependable  information  available,  using 
the  searching  light  of  science  to  wipe  out  prejudices,  half  truths, 
and  incorrect  beliefs.  The  methods  and  knowledge  of  science  are  in 
our  hands.  We  can  live  more  fully,  more  satisfyingly,  more  com- 
fortably, more  humanely,  and  more  intelligently  only  by  using  them. 
Look  forward  to  the  future,  with  the  assurance  that,  through  the 
j^fforts  of  each  of  us,  science  amafld  w  ill  serve  mahkTncT 

What  is  the  world  made  of?  For  thousands  of  years,  men  have 
tried  to  find  out  what  the  material  things  of  the  world  are  made  of. 
Socrates  (sok'rd-tez) ,  a  learned  Greek  teacher  who  lived  in  the  fifth 
century  B.C.,  believed  that  he  might  discover  the  answer  to  this  ques- 
tion simply  by  thinking  about  it.  One  other  Greek  teacher  is  said 
to  have  put  out  his  eyes  so  that  his  thinking  might  not  be  disturbed 
or  influenced  by  what  he  could  see  about  him.  He  wanted  to  give 
himself  over  completely  to  undisturbed  thinking,  or  contemplation. 

To  be  sure  these  were  extreme  cases.  Not  all  the  ancient  teachers 
relied  wholly  on  pure  thought  as  they  tried  to  find  out  what  the 
world  is  made  of.  Some  of  them,  of  course,  were  influenced  by  what 
they  could  observe.  For  example,  Thales  (tha'lez) ,  who  was  born  in 
Miletus  in  Asia  Minor  in  624  B.C.,  noticed  that  water  nourishes  crops 
and  that  it  is  found  in  large  amounts  in  the  bodies  of  men  and  other 
living  things.  Hence  he  speculated,  or  guessed,  that  water  was  the 
basic  substance  from  which  all  material  things  were  made. 

The  speculations  of  Thales  were  more  important  to  him  than  his 
observations.  It  is  said  that  one  day,  as  he  was  walking  along  ab- 
sorbed in  deep  thought  and  looking  up  at  the  sky,  he  fell  into  a  well. 
Thratte,  a  housemaid,  saw  the  accident  and  laughingly  said:  "In  his 
zeal  for  things  in  the  sky,  he  does  not  see  what  is  at  his  feet."  No 
doubt,  most  of  you  know  how  easy  it  is  to  become  so  absorbed  in 
trying  to  solve  a  problem  that  you  become  "blind"  to  common,  every- 
day things  and  may  even  be  called  "absent-minded."  In  a  similar 
manner,  it  is  possible  to  become  so  engrossed  in  thinking  about  a 
problem  that  simple  methods  of  solving  it  do  not  occur  to  the 
thinker.  Some  of  the  ancient  teachers  were  victims  of  this  "blindness." 

Careful  experiments  and 
constant  observation  are 
essential  to  the  chemist's 
study  of  matter. 

Many  of  Thales'  conclusions  concerning  what  the  world  is  made 
of  were  wrong.  He  and  other  scholars  like  him  depended  too  much 
upon  what  is  now  called  abstract  thought,  and  altogether  too  little 
upon  careful  experimentation  and  observation.  II  they  had  depended 
less  upon  abstract  thought  and  more  upon  observation  and  experi- 
mentation, accurate  information  would  have  accumulated  much 
more  rapidly.  However,  even  though  many  of  their  conclusions  were 
wrong,  their  work  was  of  value.  They  were  thinking  about  the  things 
that  surrounded  them,  which  was  more  than  was  being  done  by  most 
of  the  people  who  lived  at  that  time. 

Of  course,  in  answering  any  question,  careful  thinking  is  neces- 
sary, but  it  must  be  thinking  based_iipon  act 'uralc  Jacls,  which  can 
comc^only  from  careful  observation.  Many  of  the  ancient  teachers 
were  fine  thinkers.  They  began  with  what  they  knew,  or  thought 
they  knew,  and  reasoned  logically  to  a  conclusion.  Many  of  their 
conclusions  were  false,  however,  because  they  based  their  thinking 
on  inaccurate  or  incomplete  information. 

Ours  is  a  complex  world.  This  age-old  search  to  find  out  what  the 
world  is  made  of  is  not  over.  Far  from  it.  You,  living  2500  years  after 
Thales,  know  that  water  is  not  the  basic  substance  from  which  the 
world  is  made.  Yet,  if  you  were  asked  to  name  all  the  different  ma- 
terials that  make  up  the  world,  you  would  doubtless  think  you  had 
been  given  an  impossible  task.  Imagine  listing  all  the  materials  that 
make  up  the  rocks  and  soil  beneath  your  feet,  the  vast  expanse  of 
water  that  covers  three-quarters  of  the  earth,  the  multitude  of  liv- 
ing things,  the  deep  envelope  of  air  that  surrounds  our  globe,  and 
the  billions  of  stars,  some  so  far  from  the  earth  that  their  distances 
stagger  our  imagination! 

At  extremely  low  tempera- 
tures, a  gas  may  become 
liquid  or  even  solid.  This 
flask  contains  liquid  helium 
and  solid  air.  Because  the 
frozen  air  at  a  temperature 
of  —340°  F.  is  much  "hot- 
ter"  than  the  liquid  helium 
at  -452°  F.,  it  is  causing 
the  helium  to  boil  and  over- 
flow  the  flask. 

Westinghouse  Electric   Corporation 

Matter  exists  in  three  states.  Even  though  you  could  not  name  all 
the  substances  that  make  up  our  world,  you  would  at  least  know  that 
some  of  them  exist  as  solids,  others  as  liquids,  and  still  others  as 
gases.  In  fact,  probably  you  know  that  mailer,  which  is  anything  that 
has  weight  and  takes  up  space,  occurs  in  one  of  three  conditions  — 
solid,  liquid,  or  gaseous.  In  which  of  these  conditions,  or  slates  of 
matter,  any  substance  exists  depends  partly  on  the  nature  of  the  sub- 
stance itself,  partly  on  its  temperature,  and  partly  on  the  pressure. 

Many  substances  exist  in  all  three  states,  depending  on  the  tem- 
perature. For  example,  water  (a  liquid)  may  be  changed  to  ice  (a 
solid)  by  cooling,  or  to  water  vapor  (a  gas)  by  heating.  Changes 
from  one  state  of  matter  to  another  by  heating  or  cooling  are  very 
common.  Iron,  which  we  know  as  a  hard  gray  solid,  is  melted  in 
foundries  and  changed  to  a  shimmering,  silvery  liquid.  If  its  tem- 
perature is  raised  high  enough,  gaseous  iron  vapor  is  boiled  oft. 
While  iron  vapor  is  something  with  which  most  of  us  are  not  fa- 
miliar, astronomers  know  that  iron  exists  normally  in  the  gaseous 
state  in  certain  extremely  hot  stars. 

What  is  a  physical  change?  When  water  changes  to  steam  or  to 
ice,  only  its  form  has  been  changed.  Steam  is  a  form  of  water;  ice  is 
also  a  form  of  water.  When  a  piece  of  limestone  is  pulverized,  only 
its  form  has  been  changed.  The  small  pieces  are  still  limestone,  al- 
though their  form  differs  from  that  of  the  huge  chunks  in  which  they 
came  from  the  quarry.  These  changes  in  the  water  and  in  the  lime- 
stone have  been  in  form  only.  Neither  the  water  nor  the  limestone 


was  changed  into  another  substance.  A  change  in  which  the  original 
substance  does  not  change  into  one  or  more  other  substances  is  a 
physical  change. 

Break  a  piece  of  wood  in  two.  Heat  a  piece  of  iron  wire  in  a  vac- 
uum by  passing  an  electric  current  through  it.  The  heat  produced 
causes  the  wire  to  glow,  but,  after  the  current  is  shut  off,  the  wire 
returns  to  its  original  condition.  In  each  case,  what  kind  of  change 
has  occurred? 

How  does  a  chemical  change  differ  from  a  physical  change?  You 
know  that  when  a  piece  of  paper  burns,  it  is  completely  and  radi- 
cally changed.  The  hot  gases  that  are  given  off  and  the  ash  that  is 
left  behind  do  not  in  any  way  resemble  the  original  substance.  When 
gasoline  is  burned  in  an  engine,  the  resulting  substances  are  entirely 
different  from  the  liquid  gasoline.  Animal  tissue  is  totally  different 
from  the  vegetable  substances  from  which  it  is  made.  The  dull  tar- 
nish on  silverware  differs  completely  from  the  gleaming  silver. 

In  these  changes  more  than  the  form  of  the  original  substance 
was  altered.  In  each  case  the  composition  of  the  original  substance 
changed.  Some  form  of  energy,  usually  heat,  was  either  liberated 
or  absorbed.  A  change  in  which  the  original  substance  disappears 
(changes]  and  new  substances  are  formed  is  a  chemical  change. 

Place  a  small  quantity  of  sugar  in  a  beaker  or  other  glass  vessel. 
On  the  sugar,  pour  a  small  quantity  of  sulfunc  acid  (a  liquid  which 
is  discussed  in  detail  in  Chapter  13) .  The  white  sugar  changes  to  a 
black  spongy  substance,  which  cannot  be  dissolved,  or  absorbed,  in 
water.  What  kind  of  change  has  occurred? 

Chemistry  is  the  science  that  deals  with  the  composition  of  matter 
and  with  the  many  chemical  changes  which  matter  undergoes. 

How  is  a  substance  identified?  Telling  one  substance  from  an- 
other is  called  identifying  a  substance.  When  you  try  to  identify  a 
substance  or  to  find  out  whether  a  substance  has  undergone  a  chemi- 
cal or  physical  change,  you  need  to  know  its  characteristics  —  what 
the  substance  is  like,  and  how  it  acts  with  other  substances. 

We  find  out  what  a  substance  is  like  by  asking  such  questions  as: 
What  is  its  color?  its  odor?  its  taste?  Is  it  a  solid?  a  liquid?  a  gas? 
At  what  temperature  does  it  boil?  At  what  temperature  does  it 
freeze?  How  hard  is  it?  Does  it  conduct  electricity?  The  answers  to 
these  questions  are  characteristics  that  enable  us  to  describe  and 
identify  a  substance.  These  characteristics  of  a  substance  are  called 
its  physical  properties. 

We  find  out  how  a  substance  acts  with  other  substances  by  placing 
it  in  contact  with  these  substances  and  observing  what  occurs.  We 


ask  also  how  sunlight,  electricity,  and  heat  affect  it.  Our  observations 
and  the  answers  to  these  and  other  questions  give  additional  charac- 
teristics that  enable  us  to  describe  and  identify  a  substance.  These 
characteristics  are  called  the  chemical  properties  of  that  substance. 

Since  two  different  substances  never  have  exactly  the  same  physical 
and  chemical  properties,  any  substance  may  be  identified  by  deter- 
mining these  properties. 

The  ancients  believed  the  world  made  of  "four  elements."  But 
to  return  to  our  original  question:  What  is  the  world  made  of? 
Guesses  and  speculations  would  be  useless  in  attempting  to  answer 
this  question.  Because  the  ancients  depended  chiefly  upon  these  pro- 
cedures and  also  upon  inaccurate  and  uncontrolled  experimentation 
and  observation,  they  made  little  progress  in  answering  the  question. 

After  Thales  had  suggested  water,  another  man  proposed  that  air 
might  be  another  of  the  basic  substances  from  which  all  matter  was 
made.  Fire,  too,  was  suggested  and  later  earth.  Pythagoras  (pi-thag'6- 
ras) ,  an  ancient  Greek  thinker  and  mathematician  who  lived  about 
600  B.C.,  is  thought  to  have  been  the  first  European  to  express  the 
idea  that  all  matter  was  composed  of  these  "four  elements." 

These  conclusions  seemed  to  be  proved  by  the  observations  of  the 
early  investigators.  When  a  stick  of  green  wood  was  burned,  they 
saw  that  lire  was  produced,  water  was  forced  out  and  boiled  off  at 
the  ends  of  the  stick,  a  smoky  vapor  (air)  was  given  off,  and  an  ash 
(earth)  was  left  behind.  They  concluded,  therefore,  that  all  matter 
was  made  up  of  different  amounts  of  two  or  more  of  these  four  basic, 
or  elementary,  substances. 

The  Greek  thinkers,  however,  made  a  serious  mistake.  They  failed 
to  make  enough  observations  of  different  substances.  They  did  not 
make  enough  experiments.  Consequently,  their  conclusions  were 
wrong.  Strangely  enough,  the  idea  that  all  matter  is  composed  of 
"four  elements"  (earth,  air,  fire,  and  water)  persisted  until  the 
eighteenth  century  and  was  considered  correct  by  many  otherwise 
well-informed  persons.  Even  today  speakers  and  writers  often  refer 
to  the  violent  actions  of  air  and  water  as  "the  fury  of  the  elements." 

The  universe  is  made  of  an  even  100  chemical  elements.  Scien- 
tists now  consider  that  the  mountains,  the  oceans,  the  air,  all  living 
things,  and  even  the  stars  and  the  rest  of  the  universe  are  composed 
of  simple  natural  substances  that  cannot  be  broken  down,  or  de- 
composed, into  simpler  substances  by  the  ordinary  types  of  chemical 
change.  A  substance  that  cannot  be  broken  down,  or  decomposed, 
into  a  simpler  substance  by  the  ordinary  types  of  chemical  change  is 
an  element. 

These  natural  elements  do  not  all  occur  on  earth  in  equal 
amounts.  Taken  together,  20  of  them  make  up  99.5  percent  of  the 
weight  of  the  crust  of  the  earth.  All  the  other  elements  comprise 
only  O.f>  percent  of  its  weight.  In  all.  there  are  100  elements.  Eight 
of  these  elements  have  been  produced  in  laboratories  by  scientists. 
Probably  the  elements  with  which  you  are  most  familiar  are  gold, 
silver,  iron,  copper,  nickel,  tin,  aluminum,  sulfur,  oxygen,  carbon, 
nitrogen,  and  hydrogen.  A  list  of  all  the  elements  is  given  on  page 

What  is  a  compound?  Most  of  the  substances  we  see,  such  as  sand, 
chalk,  cotton,  table  salt,  and  water,  are  not  elements.  Rather,  each 
is  composed  of  two  or  more  elements  so  combined  that  (1)  only 
chemical  action  can  tear  them  apart,  and  (2)  the  elements  of  which 
each  substance  is  composed  can  no  longer  be  identified  by  their 
original  individual  properties.  A  substance  composed  of  two  or  more 
elements  .\o  combined  that  the  elements  can  no  longer  be  identified 
by  their  original  individual  properties  is  a  compound.  The  elements 
of  which  a  compound  is  composed  are  said  to  be  chemically  com- 
bined, or  themically  united. 

Marble,  for  example,  is  a  compound  made  up  of  three  elements, 
carbon,  calcium,  and  oxygen,  chemically  combined.  The  properties 
of  a  compound,  such  as  color,  odor,  taste,  form,  and  ability  to  dis- 
solve in  water,  are  nearly  always  distinctly  different  from  the  prop 
erties  of  the  elements  of  which  it  is  composed.  For  example,  pure 
cane  sugar,  a  sweet,  white,  crystalline  solid  which  dissolves  in  water, 
is  completely  different  from  any  and  all  of  the  three  elements  of 
which  it  is  composed. 

How  does  a  mixture  differ  from  a  compound?  In  a  compound 
the  elements  must  be  chemically  united.  But  there  are  other  kinds 
of  substances  made  up  of  two  or  more  elements  or  compounds.  Al- 

or  mixed,  each  of  the  original  substances  can  still  be  identified  by 
its  original  individual  properties.  Hence,  the  substances  are  not 
chemically  united. 

A  pinch  of  salt  and  a  pinch  of  white  sand  stirred  together  make  an 
excellent  example  <>!'  one  of  these  substances.  The  salt  can  be  iden- 
tified by  its  characteristic  taste  and  the  sand  by  its  gritty  feel  on  the 
tongue  and  teeth.  A  substance  composed  of  two  or  more  elements  or 
compounds  that  still  retain  their  individual  properties  after  they 
have  been  thoroughly  mixed  is  a  mixture.  Some1  of  the  most  useful 
substances  in  the  world,  such  as  soil,  air,  petroleum,  and  milk  and 
many  other  foods,  are  mixtures. 

The  properties  of  a  mixture  are  the  same  as  the  properties  of  the 
elements  or  compounds  that  compose  it.  A  handful  of  iron  powder 
mixed  with  a  handful  of  powdered  sulfur  makes  a  mixture  that  re- 
sembles both  the  black  iron  and  the  yellow  sulfur.  If  a  magnet  is 
passed  through  it,  the  iron  clings  to  the  magnet.  If  a  liquid  (ailed 
carbon  disulfide  is  added  to  the  mixture,  the  sulfur  is  dissolved.  But 
if  the  mixture  of  sulfur  and  iron  is  heated,  these  two  elements  com- 
bine, forming  a  compound  known  as  iron  sulfide.  Iron  sulfide  does 
not  look  like  either  sulfur  or  iron.  It  is  not  magnetic  and  does  not 
dissolve  in  carbon  disulfide.  The  properties  of  this  compound  do  not 
resemble  those  of  either  sulfur  or  iron. 

Substances  in  certain  mixtures  may  be  separated  mechanically.  A 
mixture  of  salt  and  sand  may  be  separated  by  adding  water.  The 
salt  dissolves  and  the  sand  settles  to  the  bottom.  A  mixture  of  iron 
and  sulfur  may  be  separated  by  passing  a  magnet  through  it  or  by 
adding  carbon  disulfide  which  dissolves  the  sulfur. 

The  phlogiston  theory,  an  erroneous  explanation  ol  burning. 
One  of  man's  greatest  early  achievements  was  the  discovery  of  the 
use  of  fire.  So  strange  did  fire  appear  that  for  a  long  time  men  wor- 
shiped it.  They  considered  it  the  force  responsible  for  all  creation. 
They  pondered  over  its  mystery  and  made  many  attempts  to  explain 
it.  Karly  alchemists  thought  that  fire  was  the  result  of  some  vague 
"sulfur"  which  burnable  substances  contain.  But  later  alchemists 
felt  the  need  for  a  better  explanation  —  an  explanation  which  took 
into  account  more  of  the  facts  that  had  been  observed  in  burning 
^  many  different  substances.  A  statement  that  takes  into  account  and 
attempts  to  explain  observed  tacts  is  known  as  a  theory. 


Over  many  centuries,  alchemists,  the  forerunners  of  modern  chemists,  worked  in 
vain  with  their  crude  equipment  to  find  the  secrets  of  prolonging  life  and  of 
making  gold  from  base  metal. 

About  300  years  ago  Becher  (bek'er)  ,  a  German  scientist,  ad- 
vanced the  theory  that  all  burnable  substances  contain  phlogiston 
(flo-jis't6n)  ,  or  "fire  stuff."  He  said  that  when  a  substance  burned, 
phlogiston  left  it  in  the  form  of  (lame.  Becher  thought  that  the  ash 
formed  when  a  substance  burned  was  the  substance  in  in  us  its  phlo- 
giston. According  to  his  theory,  substances  that  burn  readily,  leaving 
little  ash.  contain  a  great  deal  of  phlogiston,  while  substances  that 
burn  with  difficulty  and  leave  much  ash  contain  little.  The  phlo- 
giston theory  was  the  first  great  theory  in  chemistry. 

The  phlogiston  theory  seemed  correct  to  the  alchemists  because  of 
certain  observations  they  had  made.  A  rising  candle  flame  seems  to 
tug  at  the  wick.  To  the  alchemists  this  suggested  that  phlogiston  was 
escaping  from  the  binning  candle.  \Vhen  a  small  amount  of  pow- 
dered lead  is  heated  in  an  iron  spoon,  it  melts,  burns,  and  forms  a 
yellow  powder.  According  to  the  phlogiston  theory,  this  yellow  pow- 
der is  lead  ash,  or  lead  minus  its  phlogiston.  Now  if  some  way  could 
be  found  to  add  phlogiston  to  this  lead  ash,  lead  should  be  produced 
again.  Perhaps  this  could  be  done  by  heating  the  lead  ash  on  some 
substance  that  contains  a  lot  of  phlogiston,  such  as  charcoal.  The 
charcoal  might  give  up  some  of  its  phlogiston  to  the  lead  ash.  When 
this  experiment  is  performed,  the  final  product  is  actually  lead. 

You  can  see  that  this  experiment  seems  to  prove  the  correctness 
of  the  phlogiston  theory.  Why?  What  mistake  did  the  early  sci- 
entists make  in  attempting  to  prove  the  phlogiston  theory? 

For  more  than  two  centuries  the  phlogiston  theory  was  considered 
to  be  an  accurate  explanation  of  burning,  and  many  of  the  famous 
pioneers  of  modern  chemistry,  among  them  Priestley  (prest'll)  and 
Schcele  (sha'l<?)  were  its  ardent  supporters. 

The  first  clue  to  the  true  explanation  of  burning.  Kven  though 
<the  phlogiston  theory  seemed  to  be  upheld  by  experiments  similar 
to  the  one  with  lead,  other  experiments  showed  it  to  be  false.  In 
these  experiments  the  fact  that  many  substances  increase  in  weight 
when  burned  could  not  be  explained  on  the  basis  of  the  phlogiston 
theory,  for  according  to  it,  all  substances  lose  phlogiston  when 
burned,  and  thereby  lose  weight. 



Lavoisier  Priestley 

In  1774  Joseph  Priestley,  an  English  minister  and  amateur  scien- 
tist, led  the  way  toward  a  true  explanation  of  burning  by  his  dis- 
covery of  a  gas  later  named  oxygen.  Because  of  political  and  reli- 
gious persecution,  he  fled  from  England  and  spent  his  last  years  in 
Northumberland,  Pennsylvania.  Priestley  showed  that  the  gas  he 
had  discovered  was  present  in  air  and  was  closely  connected  with 
burning.  Let  us  trace  the  steps  that  led  to  the  discovery  of  the  true 
nature  of  burning. 

1)  After  his  discovery  of  oxygen,  Priestley,  like  a  true  scientist, 
did  not  keep  his  discovery  to  himself.  While  he  was  in  Paris  later 
on  in  1774,  he  visited  Lavoisier    (la-vwa-zya') ,  the  most  eminent 
chemist  in  France,  and  told  him  about  his  discovery.  Priestley's  in- 
formation was  a  welcome  addition  to  the  many  facts  which  Lavoi- 
sier had  already  collected  regarding  the  nature  of  burning. 

2)  Lavoisier  examined  carefully  all  the  facts  that  he  knew  about 
burning.  He  pondered  over  them  for  months,  trying  to  formulate 
an  accurate  theory  that  would  explain  burning  and  be  in  keeping 
with  all  the  observed  facts. 

3)  Lavoisier  was  genius  enough  to  use  Priestley's  work  as  the 
basis  of  a  theory  which  would  explain  the  age-old  puzzle  of  burning. 
He  suspected,  and  later  advanced  the  theory,  that  when  a  substance 
burned,  it  increased  in  weight  because  it  united  with  something  that 
was  present  in  air.  Later  this  something  was  shown  to  be  oxygen. 

4)  Other  scientists  had  noticed  this  increase  in  weight  when  sub- 
stances were  burned  in  air.  However,  Lavoisier  was  the  first  to 
formulate  a  theory  based  on  this  fact.  He  also  undertook  a  series 
of  careful  experiments  to  see  whether  their  results  would  prove  his 
theory  to  be  correct. 


Lavoisier's  classic  12-day  experiment  which  explained  burning. 

"I  introduced  four  ounces  of  pure  mercury  into  a  [sealed]  glass  ves- 
sel," he  wrote.  "I  lighted  a  fire  in  the  furnace  which  I  kept  up  con- 
tinually for  twelve  days.  On  the  second  day,  small  red  particles 
already  had  begun  to  appear  on  the  surface  of  the  mercury."  When 
most  of  the  mercury  had  been  converted  into  a  red  powder,  he  re- 
moved the  glass  vessel  and  its  contents  (which  he  had  weighed  before 
the  experiment)  and  weighed  them  again.  There  was  no  increase 
in  weight. 

Since  the  glass  vessel  was  sealed,  nothing  had  entered  or  escaped 
from  it  during  the  heating.  Yet  when  he  broke  the  seal  he  noticed 
that  air  rushed  into  the  vessel.  To  him  this  inrush  of  air  indicated 
that  part  of  the  air  in  the  vessel  had  been  used  up  during  the  heat- 
ing, and  had  left  space  for  more  air  to  enter.  After  air  had  entered 
the  vessel,  he  weighed  it  once  more  and  determined  the  increase  in 
weight.  He  concluded  that  this  increase  in  weight  equaled  the  weight 
of  something  in  the  air  in  the  vessel  that  must  have  combined  with 
the  mercury,  forming  the  red  powder. 

Lavoisier's  inquiring  spirit  was  not  satisfied.  He  was  a  scientist  in 
the  most  modern  sense.  He  refused  to  jump  to  a  hasty  conclusion  on 
the  basis  of  a  single  experiment.  He  withheld  drawing  a  conclusion 
until  he  had  performed  many  more  experiments.  As  a  further  pre- 
caution, he  reversed  his  original  experiment.  He  took  the  red  pow- 
der of  mercury  and  heated  it  to  a  higher  temperature.  He  found 
that  all  of  the  red  powder  was  changed  back  into  mercury  and  that 
a  gas  was  given  off,  which  he  found  by  a  series  of  tests  to  be  identi- 
cal with  the  oxygen  gas  that  Priestley  had  discovered.  Hence,  he  con- 
cluded that  it  was  the  oxygen  in  the  air  that  was  responsible  for 
burning.  Of  all  the  substances  he  tried,  he  found  none  that  could 
burn  without  oxygen. 

5)  Burning,  said  Lavoisier,  is  the  chemical  union  of  a  burnable 
substance  with  oxygen.  Simple  enough.  No  mysterious  phlogiston, 
and  the  testimony  of  the  most  sensitive  balance  in  Europe  to  sup- 
port his  reasoning.  Thus,  Lavoisier  discovered  the  true  explanation 
of  burning. 

6)  Lavoisier  repeated  the  original  experiment  using  other  sub- 
stances, including  tin  and  sulfur.  He  found  that  the  results  of  these 
experiments  were  fully  in  accord  with  his  theory.  In  this  way  his 
theory  was  given  further  support. 


"Nothing  happens  without  a  cause/9  said  Leucippus  2500  years 
ago.  Ever  since  man  first  appeared  on  earth,  he  has  been  working 
constantly  to  find  out  the  why  of  many  natural  occurrences. 

What  are  some  of  these  occurrences  that  man  has  tried  to  explain? 
Although  these  are  only  a  few,  we  might  include:  Why  does  a  stone 
fall?  What  is  fire?  Why  do  some  substances  burn  while  other  sub- 
stances do  not?  What  makes  thunder?  Why  can  birds  fly? 

In  asking  ourselves  these  questions  and  in  answering  them,  we  use 
several  words  that  probably  were  not  used  by  primitive  man.  These 
words,  which  appear  later  on  in  this  paragraph,  are  words  whose 
meanings  have  come  into  rather  common  use  comparatively  recently, 
as  scientists  measure  time,  perhaps  within  the  last  25,000  or  50,000 
years.  Let  us  examine  the  first  question.  Probably  we  would  ask: 
What  causes  a  stone  to  fall?  By  this  we  mean:  What  force  causes  a 
stone  to  fall,  for  we  know  that  a  stone  will  not  fall  unless  some  lorce 
acts  on  it,  producing,  we  might  say,  an  effect.  We  would  answer  the 
question:  A  stone  falls  because  the  earth  pulls  the  stone  toward  its 
center  (the  force  of  gravity).  As  Leucippus  (lu-slp'iis)  implied, 
modern  scientists  and  most  modern  people  believe  that  every  cause 
has  an  effect,  and  every  effect  has  a  cause.  Such  a  relationship  is 
known  as  a  cause-and-effect  relationship. 

Establishing  a  cause-and-effect  relationship  is  not  as  simple  as  it 
might  seem.  Early  man  did  not  have  the  many  tools  and  instruments 
which  today  we  use  so  casually  in  finding  cause-and-effect  relation- 
ships. Consequently,  in  attempting  to  explain  the  causes  of  a  certain 
effect,  early  man  relied  on  what  we  would  now  consider  magic,  mys- 
ticism, and  superstition,  but  later  on  man  learned  to  establish  these 
relationships  by  other  methods.  Then,  too,  it  is  sometimes  difficult 
to  establish  a  cause,  because  often  several  causes  taken  together  pro- 
duce a  single  simple  effect.  Today  many  cause-and-effect  relation- 
ships are  clearly  understood;  but  on  the  other  hand,  the  causes  of 
certain  effects  are  not  yet  known,  or  even  when  known,  are  not  thor- 
oughly understood. 

The  method  of  deduction  compared  with  induction.  In  explain- 
ing a  natural  occurrence,  Aristotle  (ar'Is-tot"l) ,  a  well-known  teacher 
and  philosopher  of  ancient  Greece,  often  made  a  bold  and  sweeping 
general  statement.  From  this  general  statement  he  drew  inferences 
and  conclusions,  which  he  thought  applied  in  other  similar  cases. 
Aristotle's  method  is  commonly  known  as  the  method  of  deduction. 

Francis  Bacon,  who  lived  almost  20  centuries  later,  was  the  first 
man  to  make  popular  another  method  of  reasoning.  By  a  process  con- 
sisting of  observation,  collection  of  facts  concerning  the  problem, 


formulation  of  a  theory  taking  into  account  and  explaining  the 
observed  facts,  and  verification  of  the  theory  by  actual  experiment, 
he  formulated  broad  principles,  sometimes  called  laws.  Bacon's 
method  is  known  as  the  method  of  induction,  and  today  it  is  used 
widely.  It  is  the  pattern  on  which  scientific  method  is  based.  To  a 
great  extent,  it  is  responsible  for  the  success  of  scientific  method. 

While  the  method  of  induction  is  used  very  widely,  deduction  has 
played  its  part  too  in  the  development  of  science.  The  establishment 
of  cause-and-effect  relationships,  usually  by  the  method  of  induction, 
is  perhaps  the  greatest  function  of  science. 

What  is  scientific  method?  The  method  that  Lavoisier  used  in 
reaching  the  first  correct  explanation  of  burning  is  an  example  of 
a  pattern  of  action  and  thought  used  by  scientists  in  their  work.  This 
pattern  is  known  as  scientific  method. 

Scientific  method  may  vary  according  to  the  nature  of  the  problem 
to  be  solved  and  the  tools  available  for  solving  it.  In  general,  how- 
ever, the  steps  of  the  scientific  method  are  represented  by  the  six 
steps  you  have  just  traced.  In  brief,  they  may  be  stated  as  follows: 

The  collection  of  all  available  facts  related  to  a  problem 

The  open  •  minded  checking  and  examination  of  these  facts 

The  formulation  of  a  working  theory  based  upon  these  facts 
The  testing  of  this  working  theory  by  experiments 

The  formulation  of  a  law,  or  principle,  from  the  tested  theory 
The  use  of  this  law  in  concrete  and  specific  situations 

tr ' '. 

As  you  see,  a  scientific  law,  or  principle,  is  a  descriptive  and  ex- 
planatory statement,  or  generalization,  that  expresses  what  men  have 
found  to  be  accurate  with  respect  to  natural  occurrences. 

Triumph  of  scientific  method.  Lavoisier's  explanation  of  burning 
was  not  at  once  accepted.  Indeed,  for  some  time  it  met  with  bitter 
opposition.  Those  who  believed  in  the  phlogiston  theory  attempted 
to  adjust  their  theory  to  fit  the  newly  discovered  facts,  but  this  could 
not  be  done.  Even  so,  Lavoisier  realized  how  hard  it  would  be  to 


convince  everyone  of  the  truth  of  his  own  theory.  He  wrote:  "I  do 
not  expect  that  my  ideas  will  be  accepted  at  once;  the  human  mind 
inclines  to  one  way  of  thinking,  and  those  who  have  looked  at  na- 
ture from  a  certain  point  of  view  during  a  part  of  their  lives  adopt 
new  ideas  only  with  difficulty;  it  is  for  time,  therefore,  to  confirm 
or  reject  the  opinions  that  I  have  advanced.  Meanwhile,  I  see  that 
young  men  who  are  beginning  to  study  the  science  (chemistry) 
without  prejudice  or  preconceived  notions  no  longer  believe  in 

New  ideas,  discoveries,  and  new  theories  as  a  rule  must  overcome 
tradition  and  prejudice,  but  this  should  not  discourage  those  who 
introduce  them.  Tradition  and  prejudice  are  found  not  only  in  the 
field  of  the  natural  sciences  —  physics,  chemistry,  astronomy,  biol- 
ogy, and  others  —  but  also  in  the  social  sciences  —  government,  eco- 
nomics, sociology,  history,  and  others  —  in  which  they  are  likely  to 
be  even  more  pronounced. 

Lavoisier's  theory  of  burning  finally  triumphed,  however.  The 
accumulated  evidence  in  its  favor  finally  became  so  overpowering 
that  scientists  could  believe  nothing  else.  The  experimental  method 
of  science  had  won  over  the  strictly  logical  and  theoretical  method  of 
the  ancients.  Twenty  years  later  Lavoisier,  who  was  an  aristocrat,  was 
beheaded  during  the  frenzy  of  the  French  Revolution.  "Until  it  is 
realized  that  the  gravest  crime  of  the  French  Revolution  was  not  the 
execution  of  the  king,  but  of  Lavoisier,  there  is  no  right  measure  of 
values,  for  Lavoisier  was  one  of  the  greatest  three  or  four  men 
France  had  produced."  This  statement,  made  by  an  eminent  French- 
man, expresses  the  judgment  of  thinking  men  the  world  over. 

There  are  no  authorities  or  dictators  in  science.  There  are  only 
those  persons  who  know  what  men  up  to  now  have  discovered;  that 
is,  there  are  experts.  In  all  of  science  there  is  no  one  who  can  say 
"This  theory  is  true."  But  there  are  many  men  who  can  say  "On  the 
basis  of  what  we  now  know,  this  theory  seems  to  be  sound."  Scien- 
tific method  is  a  democratic  method.  In  most  cases,  its  theories  are 
the  outcome  of  a  pooling  of  facts  discovered  by  many  workers,  a 
cooperative  effort. 

Careful  weighing  and  the  law  of  the  conservation  of  matter.  To 
Lavoisier  a  balance  was  absolutely  necessary.  By  carefully  weighing 
all  the  substances  entering  into  his  experiments  both  before  and 
after  each  experiment,  he  found  that  there  was  no  loss  of  weight 
during  the  burning.  The  mercury  plus  the  oxygen  from  the  air 
weighed  exactly  the  same  as  the  red  powder  of  mercury  (mercuric 
oxide) .  "One  may  take  it  for  granted,"  he  wrote,  "that  in  every 



change  there  is  an  equal  quantity  of  matter  before  and  alter  the  op- 
eration." In  chemical  changes  we  can  change  the  form,  the  state,  or 
the  composition  of  matter,  but  we  cannot  destroy  matter  itself.  In 
physical  changes  we  can  change  the  ionn  or  the  state  of  matter,  but 
we  cannot  destroy  the  matter  itself. 

"But/*  you  might  answer,  "a  candle  burns  until  i(  is  all  gone.  It 
becomes  smaller  and  smaller  and  certainly  weighs  less  and  less." 
And  you  are  right.  But  if  we  take  the  trouble  to  collect  and  weigh 
all  the  gases  formed  during  the  burning  of  the  candle,  we  lind  that 
they  weigh  more  than  the  original  candle.  This  increase  in  weight 
is  due  to  the  oxygen  with  which  the  burning  candle  has  combined. 

"What  of  an  oak  tree,"  you  might  say.  "It  grows  from  a  little 
acorn.  Isn't  matter  created  here?"  It  might  seem  so,  but  the  fact  that 
many  things  increase  gradually  in  size  and  weight  does  not  mean 
that  matter  has  been  or  is  being  created.  The  oak  tree  does  not  come 
from  seed  alone.  The  cells  from  which  the  tree  is  made  are  built  up 
chiefly  from  food  materials  taken  out  of  the  air,  water,  and  soil. 
Parts  of  the  air,  water,  and  soil  have  been  chemically  changed  and 
combined,  forming  living  substances. 

Matter  can  be  neither  created  nor  destroyed.  This  fundamental 
law  both  of  chemistry  and  of  all  science,  is  called  the  law  of  the  con- 
servation of  matter. 

1.  Th«  a*  o 

III*  At 

th«  ttt« 

th*  the 

tilti  on  ttt» 





Energy,  too,  can  be  neither  created  nor  destroyed.  So  far  we  have 
considered  mainly  the  changes  in  matter  that  take  place  in  burning. 
But  there  are  other  changes  that  are  equally  important.  When  a 
substance  burns,  heat  is  liberated.  This  heat  may  be  used  to  convert 
water  in  a  boiler  into  steam.  The  pressure  of  this  steam  may  then  be 
used  to  turn  a  wheel,  thus  producing  rotary  motion.  By  connecting 
this  wheel  to  a  dynamo,  electricity  may  be  generated.  This  electricity 
may  be  changed  into  heat  or  light  or  magnetism,  depending  on 
whether  it  is  sent  through  a  toaster,  a  light  bulb,  or  the  electromag- 
net of,  let  us  say,  a  buzzer  or  telegraph  sounder.  Or  this  electricity 
may  be  used  to  charge  a  storage  battery.  In  charging  the  battery,  the 
electricity  produces  a  chemical  change,  which,  on  being  reversed, 
yields  electricity  again. 

Evidently  all  these  —  heat,  electricity,  the  power  to  produce  mo 
tion,  light,  magnetism,  and  the  power  to  produce  a  chemical  change 
—  can  be  transformed  one  into  the  other.  All  are  capable  of  doing 
work,  and  all  are  forms  of  energy.  In  all  energy  changes,  just  as  in 
all  changes  in  matter,  there  is  no  loss,  only  transformation.  Energy 
can  be  neither  created  nor  destroyed.  This  is  a  fundamental  law  both 
of  chemistry  and  of  all  science.  It  is  called  the  law  of  the  conserva- 
tion of  energy. 

Two  laws  or  one?  Researches  on  the  structure  of  matter  and  the 
nature  of  energy  resulted  in  the  atomic  bomb.  These  researches, 

one  of 

be  of  energy,  follow 

in  the 

discussed  in  Chapter  12,  lead  definitely  to  the  conclusion  that  mat- 
ter and  energy  are  but  different  forms  of  the  same  thing,  and  that 
matter  can  be  converted  into  energy  and  energy  into  matter. 

As  a  result,  the  law  of  the  conservation  of  matter  and  the  law  of 
the  conservation  of  energy  are  no  longer  considered  separate  and 
distinct  laws.  Instead,  they  may  be  considered  as  different  phases  of 
a  single  law.  Such  a  law  would  state  that  matter  and  energy  can  he 
neither  created  nor  destroyed,  but  that  each  can  be  transformed 
into  the  other.  In  the  transformation  of  matter  into  energy,  matter 
disappears  and  becomes  energy.  In  the  transformation  of  energy  into 
matter,  energy  disappears  and  becomes  matter. 

Can  any  form  of  energy  produce  a  chemical  change?  When  pa- 
per is  heated  to  its  kindling  temperature  (see  page  28)  ,  it  burns. 
Heat,  one  form  of  energy,  produces  a  chemical  change.  When  light, 
another  form  of  energy,  strikes  a  photographic  film,  it  causes  a  chemi- 
cal change  in  the  substances  that  coat  the  film.  Thus,  light  also  can 
cause  a  chemical  change.  When  an  electric  current  is  passed  through 
water,  it  splits  the  water  into  two  gases,  neither  of  which  resembles 
water  vapor.  Thus  electricity,  too,  is  very  effective  in  producing  a 
chemical  change.  (The  energy  that  was  used  to  split  the  water  ap- 
pears again  as  heat  energy  when  the  two  gases  are  recombined,  form- 
ing water.)  From  these  experiments  and  from  others,  we  know  that 
many  forms  of  energy  bring  about  chemical  changes. 

"  '*' "*•'- """~ 







The  literature  of  chemistry  is  filled  with  romance.  In  it  you 
Will  find  stories  of  human  struggle  and  achievement,  stories 
whose  truth  makes  them  all  the  more  worthy  to  be  read  and 
remembered.  Even  the  most  interesting  and  significant  parts 
of  this  literature  cannot  all  be  listed  in  this  book.  Neverthe- 
less, they  are  available,  and  it  will  be  well  worth  your  while 
to  read  as  many  of  them  as  you  can.  The  following  is  a  selected 
list  of  books  that  deal  with  some  of  the  topics  in  this  chapter. 

Becker,  Carl;  Painter,  Sidney;  and  Han,  Yu-Shan.  The  Past 
that  Lives  Today,  pp.  22-86.  Silver  Burdett  Co.,  New  York, 
1952.  A  fascinating  story  of  early  man,  ancient  civilizations 
and  the  science  of  the  ancients  is  told  here. 

Fabre,  Jean  H.  The  Wonder  Book  of  Chemistry,  pp.  6-69. 
Albert  &  Charles  Boni,  New  York,  1922.  A  delightful  account 
of  elements,  mixtures,  and  compounds. 

Jaffe,  Bernard.  Crucibles:  The.Story  of  Chemistry,  pp.  34-50. 
Simon  and  Schuster,  New  York,  1948.  A  simple  account  of  the 
phlogiston  theory. 

Somerville,  John.  The  Way  of  Science:  Its  Growth  and 
Method,  pp.  93-113.  Henry  Schmnan,  New  York,  1953.  A  very 
simple  illustration  ot  the  steps  ot  scientific  method. 


1.  The  three  states  of  matter  are  solid,  liquid,  and  gaseous. 

2.  In  a  physical  change,  the  original  substance  does  not 
change  into  one  or  more  other  substances. 

3.  In  a  chemical  change,  the  original  substances  disappear 
and  new  substances  are  formed. 

4.  An  element  is  a  substance  that  cannot  be  broken  down, 
or  decomposed,  into  a  simpler  substance  by  the  ordinary  types 
of  chemical  change. 

5.  A  chemical  compound  is  a  substance  composed  of  two 
or  more  elements  so  combined  that  the  elements  can  no  longer 
be  identified  by  their  original  individual  properties. 

6.  A  mixture  is  a  substance  composed  of  two  or  more  ele- 
ments or  compounds  that  still  retain  their  individual  prop- 
erties after  they  have  been  thoroughly  mixed. 

7.  Burning  is  the  chemical  change  in  which  a  burnable 
substance  unites  with  oxygen. 

8.  The  law  of  the  conservation  of  matter  states  that  mat- 
ter can  be  neither  created  nor  destroyed.  But  matter  may  be 
changed  from  one  form  to  another. 


9.  The  law  of  the  conservation  of  energy  states  that  energy 
can  be  neither  created  nor  destroyed.  But  energy  may  be 
changed  from  one  form  to  another. 

10.  Recent  researches  prove  that  matter  can  be  transformed 
into  energy,  and  that  energy  can  be  transformed  into  matter. 

11.  Establishing  accurate   cause-and-effect  relationships   is 
perhaps  the  greatest  function  of  science. 

12.  The  steps  in  the  scientific  method  include  (1)  the  col- 
lection of  all  available  facts  related  to  a  problem,    (2)  the 
open-minded  examination  of  these  facts,   (3)  the  formulation 
of  a  working  theory  based  upon  these  facts,  (4)  the  testing  of 
this  working  theory  by  experiments,  (5)  the  formulation  of  a 
law,  or  principle,  from  the  tested  theory,  and   (6)  the  use  of 
the  law  in  specific  situations. 

13.  Blind  acceptance  of  so-called  "authorities,"  prejudices, 
and  personal  likes  and  dislikes  have  no  place  in  the  general 
pattern  of  action  and  thought  of  the  true  scientist. 


Group  A 

1.  (a)  Which  of  the  following  are  physical  changes  and 
which    are    chemical    changes?     (b)  Give    reasons    for    your 
answers.  Souring  of  milk,  molding  of  clay,  digestion  of  food, 
drying  of  clothes,  dissolving  sugar  in  water,  decay  of  fruit, 
freezing  of  mercury,  photosynthesis  (manufacture  of  starch  in 
the  leaves  of  plants  from  carbon  dioxide  gas  and  water)  , 
erosion  (the  breaking  up  of  rocks  and  soil  by  the  action  of  air 
and  water)  . 

2.  (a)  What  were  the  "four  elements"  of  the  ancients? 
(b)  In  what  respects  did  their  observations  seem  to  support 
this  theory?   (c)  What  was  one  weakness  of  the  way  in  which 
this  theory  was  formulated? 

3.  What  are  your  reasons  for  thinking  that  water  is  a 

4.  Classify  the  following  as  elements,  compounds,  or  mix- 
tures: table  salt,  mercury,  aluminum,  paper,  carbon  dioxide, 
gold,  silver,  iron  rust,  sugar,  sulfur,  milk,  brass,  silver  coin. 

5.  Name  a  compound  that  contains  hydrogen. 
4.  Explain  one  method  of  telling  a  mixture  from  a  com- 
pound. «• 


7.  Give  briefly  the  main  ideas  of  the  "phlogiston  theory." 

8.  What  was  Priestley's  part  in  discovering  the  true  nature 
of  burning? 

9.  (a)  Complete  the  following  statement:  When  red  mer- 
curic oxide  is  heated,  it  is  changed  into and (b)  Is 

the  change  a  physical  or  a  chemical  change?  (c)  Why? 

I .  T . 

'  i 

10.  Describe  Lavoisier's  12-day  experiment  on  burning. 

11.  Is  Lavoisier's  explanation  the  modern  explanation  of 

12.  What  part  did  the  balance  play  in  the  development  of 

13.  State  the  six  general  steps  in  the  scientific  method. 

14.  Distinguish    between    the   method   of  science   and   the 
method  ol  the  ancient  teachers  and  philosophers. 

I  ... 

15.  (a)  What   is   the   law  of   the  conservation   of  matter? 
(b)    Is  the  disappearance  of  camphor  balls  in  clothes  an  excep- 
tion to  this  law?    (c)  Explain. 

16.  (a)  What  is  energy?   (b)  Name  three  different  forms  of 

17.  Give  an  illustration  to  show  that  each  of  the  forms  of 
eneigy  named  in  your  answer  to  question  16  may  produce  a 
chemical  change. 

18.  Assume  that  coal  is  the  source  of  energy  that  lights  your 
home    (electric   light).  Make  a  list  of  the  transformations  of 
energy  that  occur,  beginning  with  the  burning  of  coal  and  end- 
ing with  the  lighted  bulb. 

Group  B 

19.  Not  many  of  the  chemical  elements  wrere  discovered  by 
Americans,    (a)  Which  elements  are  these?    (b)  Can  you  sug- 
gest a  reason  why  Americans  have  discovered  so  few? 

20.  Aluminum  is  the  most  abundant  metal  in  the  earth. 
Tell  why  only  in  recent  years  aluminum  has  come  into  com- 
mon use. 

21.  (a)  Can  one  use  scientific  methods  in  fields  other  than 
science?  (b)  Explain  your  answer. 

22.  (a)  State  four  evidences  of  chemical  action,  and  (b)  give 
one  example  of  each. 



1.  Discuss  with  your   teacher  of  history  the  problem  of 
prejudice  and  tradition  as  obstacles  to  progress.  Write  a  report 
on  this  topic  using  an  illustration  from  American  history  or 
from  the  history  of  science. 

2.  Find  an  article  on  some  subject  such  as  mental  telepathy, 
communications  from  the  spirit  world,  extrasensory  percep- 
tion, astrology.  Read  the  article  carefully  and  write  your  own 
reaction  as  to  whether  you,  using  the  methods  of  science, 
would  accept  the  author's  conclusions  as  scientific. 

3.  Obtain  a  new  photoflash  bulb  and  weigh  it  on  a  sensitive 
balance.  Weigh  it  again  after  it  has  been  ignited.  Do  your 
findings  uphold  the  law  of  the  conservation  ol  matter? 

4.  Show  how  you  would  use  the  method  of  science  in  solving 
some  particular  everyday  problem  which  you  have  had  to 
solve  or  which  you  will  soon  have  to  solve. 

5.  With  the  aid  of  a  medicine  dropper,  allow  a  drop  of  ink 
to  fall  into  a  tall  glass  of  cold  water.  Observe  what  happens 
after  approximately  5  minutes,  8  hours,  2-1  hours,  and  2  days. 
Report    to    class    with    reference    to    physical    and    chemical 
changes,  and  any  other  conclusion  you  have  drawn. 

6.  Prepare  a  brief  report  on  an  expert  in  some  field  of  sci- 
ence. Find  out  how  he  became  an  expert.  How  does  an  expert 
differ  from  an  authority? 

7.  On  the  first  page  of  this  chapter,  there  is  a  statement 
made  by  Lavoisier.  Read  it  carefully.  Report  to  your  class  on 
the  importance  of  this  statement.  What  are  its  implications 
for  everyday  living? 


.  .  .  /  procured  a  mouse  and  put  it 
into  a  glass  vessel  containing  the  air 
(oxygen)  from  the  red  powder  of 
mercury.  Had  it  been  common  air,  a 
full-grown  mouse,  as  this  was,  would 
have  lived  in  it  about  a  quarter  of  an 
hour.  In  this  air,  however,  my  mouse 
lived  a  full  half  hour. 

Joseph  Priestley,  1775 

A  Sunday  experiment  by  an  English  minister.  On  Sunday,  the 
first  of  August,  1774,  Priestley  was  working  in  his  laboratory.  He 
placed  a  red  powder  (mercuric  oxide)  in  a  bell  jar  so  arranged  that 
any  gas  which  might  be  formed  would  pass  out  of  the  bell  jar 
through  a  tube  and  be  collected  in  a  bottle. 

Instead  of  heating  the  powder  over  a  flame,  he  used  a  large  burn- 
ing lens  to  concentrate  the  rays  of  the  sun  on  the  powder.  "I  pres- 
ently found/'  lie  reported,  "that  air  was  expelled  from  it  readily." 
But  this  result  was  not  startling,  because  others  before  him  had  ob- 
tained gases  by  heating  solids. 

A  candle  was  burning  in  the  laboratory.  Wondering  what  effect 
this  gas  would  have  on  a  flame,  he  placed  the  candle  in  a  bottle  of 
it.  Priestley  reported  in  somewhat  flowery  words:  "A  candle  burned 
in  this  air  with  an  amazing  strength  of  flame;  and  a  bit  of  red-hot 
wood  crackled  and  burned  with  a  prodigious  rapidity  exhibiting  an 
appearance  something  like  that  of  iron  glowing  with  a  white  heat 
and  throwing  out  sparks  in  all  directions.'* 

Only  natural  curiosity  or  perhaps  chance  led  Priestley  to  experi- 
ment with  the  gas  (or  as  he  called  it,  air) .  As  Priestley,  himself,  said 
later,  he  had  no  idea  what  the  outcome  might  be. 


Priestley  was  unable  to  explain  what  had  happened.  He  was  such 
a  firm  believer  in  the  phlogiston  theory  that  he  did  not  associate  this 
new  gas  with  burning.  As  we  learned  in  Chapter  1,  it  was  Lavoisier 
who  showed  that  Priestley's  air  (later  called  oxygen)  is  the  element 
necessary  for  burning,  thus  solving  a  mystery  that  had  baffled  scien- 
tists for  centuries. 

Half  the  earth  is  oxygen!  Priestley's  discovery  of  oxygen  was  a 
turning  point  in  the  development  of  chemistry.  It  is  one  of  the 
strangest  facts  in  history  that  this  element,  which  surrounds  us  every- 
where and  without  which  life  is  impossible,  was  not  obtained  pure 
until  about  180  years  ago.  This  fact  is  even  more  surprising  when 
we  realize  that  this  one  element,  oxygen,  is  present  on  earth  in  quan- 
tities equal  to  the  weight  of  all  the  other  chemical  elements  put 
together.  Sand  and  half  of  the  different  kinds  of  rocks  on  the  earth 
are  compounds  of  oxygen.  Water  contains  almost  99  percent  oxygen 
by  weight,  and  air  contains  about  21  percent  oxygen  by  volume. 

How  oxygen  is  prepared  in  the  laboratory.  Because  mercuric 
oxide,  from  which  Priestley  prepared  oxygen,  costs  about  one  dol- 
lar a  pound,  it  is  too  expensive  to  use  in  the  laboratory.  Instead,  we 
obtain  oxygen  from  a  white  crystalline  solid  called  potassium  chlo- 
rate. This  compound  is  composed  of  three  elements  —  potassium, 
chlorine,  and  oxygen.  By  applying  heat,  the  oxygen  can  be  torn 
away  and  liberated  as  free  oxygen. 

How  to  express  the  change  by  which  oxygen  is  prepared.  As  po- 
tassium chlorate  is  heated,  it  yields  potassium  chloride  plus  oxygen. 
This  is  a  chemical  reaction  which  produces  the  chemical  change  rep- 
resented by  the  equation: 

Potassium  (K)  \  -\ 

^, ,     .       /X«;  potassium 

Chlorine  (Cl)    \  -*>     rul     .          >  -f     oxygen 

,lv       I  chlorine    J  75 

Oxygen  (O)      J  J 

2KC1O3  -»     2KC1  +     302 

potassium  chlorate     — >     potassium  chloride     -f     oxygen 

The  chemical  shorthand  used  to  express  the  change  will  be  fully 
explained  later.  The  forms  of  energy  that  take  part  in  the  change  are, 
as  a  rule,  omitted  from  the  equation.  A  chemical  change  is  the  re- 
sult of  a  chemical  reaction. 

How  the  speed  of  this  reaction  can  be  increased.  The  method  of 
preparing  oxygen  just  described  has  one  serious  drawback.  Unless 
a  very  high  temperature  is  reached,  oxygen  is  liberated  very  slowly. 
Someone  discovered  that  if  a  small  amount  of  powdered  manganese 
dioxide,  a  black  solid,  is  added  to  the  potassium  chlorate  before 


heating,  oxygen  is  liberated  more  quickly  and  at  a  lower  tempera- 

At  the  end  of  the  chemical  reaction,  that  is,  when  oxygen  is  no 
longer  given  off,  the  same  amount  of  manganese  dioxide  with  which 
the  experiment  started  remains.  The  weight  of  manganese  dioxide 
has  not  been  changed  in  any  way.  None  of  its  oxygen  has  been  lib- 
erated. Since  manganese  dioxide  remains  unchanged  at  the  end  of 
the  reaction,  it  is  not  included  in  the  equation. 

A  catalyst  changes  the  speed  of  a  chemical  reaction.  Chemists 
have  discovered  that  many  chemical  reactions  can  be  speeded  up  or 
slowed  down  by  placing  a  small  quantity  of  certain  substances  in 
contact  with  the  reacting  materials.  A  substance  that  changes  the 
speed  of  a  chemical  reaction  is  called  a  catalyst,  or  a  catalytic  agent. 
The  catalyst  itself  may  undergo  some  temporary  change,  but  at  the 
end  of  the  reaction  is  present  in  the  same  state  and  quantity  as  at 
the  beginning.  Manganese  dioxide,  as  used  in  the  laboratory  prepa- 
ration of  oxygen,  is  a  catalyst.  However,  manganese  dioxide  does 
not  always  act  as  a  catalytic  agent.  It  (see  page  137)  actually  enters 
into  certain  reactions  in  which  its  composition  is  changed  perma- 

We  know  little  about  the  reasons  that  catalysts  act  as  they  do. 
We  do  know,  however,  that  many  vitally  important  chemical  re- 
actions take  place  too  slowly,  and  hence  uneconomically,  except  in 
the  presence  of  certain  catalysts.  Research  on  the  nature  of  catalysis 
is  now  being  carried  on  in  laboratories  throughout  the  world,  and  we 
will  know  much  more  about  it  before  very  long. 

The  presence  of  a  catalyst  in  a  reaction  is  sometimes  indicated  by 
writing  the  catalyst  over  the  arrow  in  the  equation.  In  the  reaction 
just  discussed,  the  presence  of  the  catalyst,  manganese  dioxide,  might 
be  indicated  thus: 

2KC1O3 >  2KC1  +  3O2 

How  oxygen  gas  is  collected.  In  preparing  oxygen,  a  mixture  of 
potassium  chlorate  and  manganese  dioxide  is  put  in  a  test  tube  and 
heated  over  a  bunsen  burner  (see  illustration  page  27) .  Connected 
to  the  test  tube  is  a  delivery  tube  which  reaches  into  a  bottle  that 
has  been  filled  with  water  and  placed  mouth  downward  in  a  pan  of 
water.  As  the  potassium  chlorate  is  heated,  it  is  broken  down,  or 
decomposed,  forming  potassium  chloride  and  oxygen. 

The  potassium  chloride  is  a  solid  and  remains  in  the  test  tube. 
The  oxygen  is  a  gas  and  passes  through  the  delivery  tube  into  the 

2.  of  if 

ii  the 

water-filled  collecting  bottle  from  which  it  displaces  water.  This 
method  of  collecting  gases  is  called  the  displacement  of  water  method, 
and  was  used  by  Priestley.  Priestley  also  collected  gases  by  the  dis- 
placement of  mercury  when  such  gases  dissolved  in  water  but  not 
in  mercury.  If  a  gas  is  collected  by  the  displacement  of  a  liquid, 
chemists  say  that  the  gas  is  collected  over  the  liquid.  Thus,  Priestley 
collected  oxygen  over  water. 

Physical  properties  of  oxygen.  In  discussing  the  physical  properties 
of  a  gas,  we  usually  consider  five  characteristics:  (1)  color,  (2)  odor, 

(3)  weight  compared  with  the  weight  of  an  equal  volume  of  air, 

(4)  the  ease  with  which  it  may  be  changed  into  a  liquid,  and   (5)   its 
absorption  by  water.  This  absorption  by  water  we  call  its  solubility 
in  water.  Just  as  solids,  such  as  sugar  and  salt,  disappear  as  solids 
when  stirred  in   water  and  are  distributed    uniformly  throughout 
the  liquid,  so  gases,  such  as  air  and  oxygen,  when  passed  through 
water,  may  likewise  be  absorbed  by  the  water.  In  the  case  of  some 
gases,  taste  is  considered  also. 

Oxygen  is  colorless  and  has  no  odor.  It  is  slightly  heavier  than 
air.  It  is  slightly  soluble  in  water.  Under  normal  conditions,  that 
is,  at  a  temperature  of  18  degrees  centigrade  and  a  pressure  of  one 
atmosphere,  or  760  millimeters  of  mercury,*  about  four  quarts  of 

*  Scientists  often  use  two  terms  to  describe  temperature  and  pressure,  normal 
conditions  and  standard  conditions.  Standard  conditions  mean  a  temperature 
of  0°  C.  and  a  pressure  of  760  mm.  of  mercury,  that  is,  one  atmosphere  (atm.) . 
Normal  conditions  mean,  in  effect,  roorn  temperature  and  the  pressure  of  1  at- 
mosphere. Throughout  this  book  a  pressure  of  760  mm.  is  assumed  when  staling 
temperatures  at  which  gases  liquefy  or  liquids  solidify.  A  millimeter  (mm.)  is 
a  small  unit  of  length  in  the  Metric  System.  Numerically  it  is  equal  to  0.001  of 
a  meter,  a  larger  unit  of  length  equal  to  39.37  inches.  For  an  explanation  of 
metric  units  and  temperature  scales,  see  pages  654  and  644  respectively.  Mcttic. 
units  of  measurement  are  used  by  scientists  in  all  countries. 


oxygen  gas  will  dissolve  in  100  quarts  of  water.  It  is  hard  to  change 
oxygen  to  a  liquid.  At  a  temperature  of  about  183  degrees  below- 
zero  centigrade  (— 183°C.)  and  under  a  pressure  of  760  millimeters, 
oxygen  is  converted  into  a  pale  blue  liquid,  which  can  be  attracted 
by  a  magnet.  At  —  219°C.  it  changes  into  a  bluish-white  solid. 

Chemically,  oxygen  is  an  active  element.  Because  oxygen  combines 
with  almost  all  other  elements,  forming  compounds  called  oxides,  it 
is  considered  to  be  very  active  chemically.  For  example,  when  iron 
is  exposed  to  oxygen,  rust,  which  is  an  oxide  of  iron,  is  formed. 

Iron  +  oxygen  — >  iron  oxide 
4Fc  +      3O2    ->    2Fc2O3 

If  iron  is  first  heated  until  it  glows  and  then  placed  in  a  bottle  of 
oxygen,  the  chemical  reaction  is  so  vigorous  that  the  iron  burns 
brilliantly,  throwing  off  sparks  of  glowing  iron  oxide.  This  sur- 
prising spectacle  actually  is  iron  burning  in  oxygen,  just  as  a  piece 
of  paper  burns  in  air! 

Differences  between  slow  and  rapid  oxidation.  When  a  sub- 
stance—  element  or  compound  —  combines  with  oxygen,  new  sub- 
stances are  formed.  This  chemical  union  of  a  substance  with  oxygen 
is  called  oxidation.  Rapid  oxidation,  such  as  the  burning  of  coal,  is 
accompanied  by  noticeable  heat  and  light.  During  slow  oxidation, 
such  as  the  rusting  of  iron  or  the  decaying  of  wood,  no  light  is  given 
off  nor  can  we  easily  detect  the  heat  because  it  is  given  off  so  slowly. 
Delicate  measurements,  however,  have  proved  beyond  doubt  that 
the  amount  of  heat  energy  liberated  is  the  same  whether  the  oxida- 
tion of  a  substance  takes  place  slowly  or  rapidly. 

The  soft  cold  light  of  the  firefly,  and  the  glow  of  some  fungi  and 
bacteria  are  caused  by  the  oxidation  of  a  complex  chemical  com- 
pound, luciferin,  which  they  produce. 

Why  some  substances  catch  fire  more  easily  than  others.  Some 
substances  catch  fire  at  low  temperatures,  but  others  require  ex- 
tremely high  temperatures  in  order  to  burn.  Every  substance  must 
be  raised  to  a  certain  definite  temperature  before  it  can  combine 
with  oxygen  at  such  a  rate  that  the  heat  produced  is  sufficient  to  keep 
the  substance  burning  without  the  addition  of  more  external  heat. 

In  starting  a  coal  fire,  we  often  begin  by  burning  paper,  which 
sets  fire  to  kindling  wood,  which  sets  fire  to  the  coal.  The  heat  given 
off  by  the  burning  paper  causes  the  wood  to  catch  fire;  the  heat  given 
off  by  the  burning  wood  in  turn  causes  the  coal  to  catch  fire.  The 
lowest  temperature  at  which  a  substance  catches  fire  and  continues 
to  burn  is  called  the  kindling  temperature  of  that  substance.  We 

Hitreau  of  Mintt,  U.S.  Depf.  of  Intu 

A  coal-dust  explosion  issues  from  the  U.S.  Bureau  of  Mines* 
Experimental  Coal  Mine  in  Pennsylvania.  Such  test  ex- 
plosions are  part  of  the  Bureau's  safety  research  program. 

say  that  a  substance  which  is  easy  to  set  on  fire  has  a  low  kindling 
temperature,  and  a  substance  which  is  difficult  to  set  on  fire  has  a 
high  kindling  temperature. 

A  substance  may  have  different  kindling  temperatures,  depending 
upon  the  size  of  the  particles  into  which  it  is  divided,  that  is,  upon 
its  state  of  subdivision.  A  solid  piece  of  iron  has  a  high  kindling 
temperature,  but  powdered  iron,  because  of  the  large  surface  that 
is  exposed  to  the  oxygen  of  the  air,  can  be  made  to  burn  readily 
in  air.  Many  dust  explosions  in  flour  mills,  starch  factories,  grain 
elevators,  and  coal  mines  are  caused  by  the  very  rapid  oxidation  of 
explosive  mixtures  of  air  and  finely  divided  materials.  A  spark 
resulting  from  static  electricity  or  friction  often  sets  off  the  explosion. 

Smut  dust  and  air  form  an  explosive  mixture  which  may  be 
ignited  by  static  electricity  during  threshing  operations.  Costly  fires 
of  this  kind  have  been  so  widespread  that  the  United  States  Depart- 
ment of  Agriculture  has  issued  a  bulletin  explaining  how  to  prevent 


The  explosion  of  a  mixture  of  coal  dust  and  air  has  been  used 
in  one  type  of  internal-combustion  engine. 

Substances  such  as  asbestos,  brick,  concrete,  and  marble  never 
catch  fire  because  they  are  already  completely  oxidized. 

Spontaneous  combustion.  Several  years  ago  there  were  widespread 
floods  in  the  Ohio  Valley.  The  lower  parts  of  thousands  of  haystacks 
in  the  Valley  were  soaked  with  water.  As  the  flood  waters  receded, 
farmers  were  pu//led  when  some  of  the  haystacks  began  to  catch  fire. 

The  explanation  of  this  phenomenon  is  simple.  During  respiration 
of  living  plant  cells,  food  materials  slowly  oxidize  and  heat  is  given 
off.  This  oxidation  is  speeded  up  by  the  presence  of  a  small  amount 
of  moisture.  The  hay  itself  also  slowly  oxidi/es,  liberating  heat.  As 
the  temperature  rises,  the  rate  of  oxidation  also  increases.  The  heat 
slowly  accumulates,  and  when  the  kindling  temperature  of  the  hay  is 
finally  reached,  it  bursts  into  flame.  Materials  catching  fire  in  this 
way  are  said  to  undergo  spontaneous  combustion.  (Combustion 
refers  to  any  chemical  reaction  which  produces  heat  and  light. 
Burning  is  only  one  kind  of  combustion.) 

Fires  have  also  been  caused  by  painters'  rags  saturated  with  linseed 
oil.  As  the  linseed  oil  slowly  oxidizes,  heat  is  given  off.  Unless  there 
is  a  free  circulation  of  air  to  carry  away  this  heat,  the  oily  rags  may 
become  hot  enough  to  catch  fire.  Thus,  you  see  why  oily  rags  should 
not  be  kept  in  poorly  ventilated  closets.  Finely  divided  coal  in  the 
closed  hold  of  a  ship,  or  in  a  poorly  ventilated  boiler  room  must  be 
sprinkled  with  plenty  of  water  from  time  to  time  to  prevent  the 
accumulation  of  heat  from  slow  oxidation. 

Phosphorus,  an  element  that  burns  spontaneously.  Ancient  alche- 
mists spent  most  of  their  time  looking  for  the  philosopher  s  stone, 
which  they  believed  would  change  lead  into  gold.  In  1669,  while 
searching  for  the  philosopher's  stone,  Hennig  Brand,  an  alchemist 
of  Hamburg,  obtained  a  new  and  strange  chemical  substance  from 
urine.  It  had  so  low  a  kindling  temperature  that  on  exposure  to 
air  it  caught  fire  immediately  and  burned,  forming  a  white  powder 
(phosphorus  oxide) .  Brand  made  a  famous  tour  of  Europe,  exhibit- 
ing this  unusual  substance.  Today  we  know  it  as  white  phosphorus, 
a  soft,  waxy  element,  now  obtained  by  a  chemical  process  from  bone 
deposits.  At  one  time  it  was  used  in  the  manufacture  of  matches 
(see  page  296) .  This  element  is  sometimes  referred  to  as  yellow 

To  keep  white  phosphorus  from  catching  fire,  it  must  be  stored 
under  water.  Do  not  touch  it  with  your  bare  fingers,  for  white 
phosphorus  will  cause  severe  burns  that  heal  very  slowly. 



Upon  exposure  to  air,  white 
phosphorus  ignites  sponta- 
neously in  a  violent  reaction. 
Note  the  protective  clothing 
worn  by  the  demonstrator. 

Monsanto  Chemical 

You  could  not  live  without  oxygen!  One  of  the  remarkable  ex- 
periments which  i'riestley  performed  showed  that  a  mouse,  placed 
in  a  bell  jar,  lived  twice  as  long  in  pure  oxygen  as  in  "common  air." 
This,  Priestley  was  unable  to  explain;  but  today  we  know  that  the 
chief  chemical  change  that  goes  on  in  the  body  of  any  living  animal 
is  slow  oxidation. 

Man  obtains  his  supply  of  oxygen  by  breathing.  When  he  inhales, 
oxygen  is  taken  into  his  lungs.  This  oxygen  passes  through  the  walls 
of  the  lungs  and  is  absorbed  by  the  red  cells  of  the  blood,  which 
carry  it  to  all  parts  of  the  body.  In  every  living  cell  or  body  dim, 
oxidation  takes  place,  liberating  heat  and  other  forms  of  energy.  This 
slow,  steady  oxidation  is  like  a  tiny  flame  which  keeps  life  going. 
Without  oxygen,  the  flame  is  snuffed  out  and  life  is  extinguished. 

With  the  exception  of  a  few  very  low  forms  of  life,  all  living 
things  take  oxygen  from  the  air.  This  oxygen  is  in  the  /rvv,  or  uiic.oin- 
bined,  state.  That  is,  the  oxygen  is  not  chemically  united  with  any 
of  the  other  substances  in  air.  Fish  obtain  their  supply  ol  oxygen, 
from  the  air  that  is  dissolved  in  water. 


Adapted  from  drawing  by  Linde  Air  Products  Company 

Fig.  3.  Oxyacetylene  torch.  What  is  the 
function  of  the  expansion  chamber? 

Industry  uses  great  quantities  of  oxygen.  Commercial  production 
of  pure  oxygen  in  the  United  States  is  more  than  25  billion  cubic 
feet  a  year.  Of  the  oxygen  produced,  it  is  estimated  that  more  than 
95  percent  is  used  in  cutting  and  processing  steel  and  in  welding 
metals,  such  as  aluminum  and  steel,  by  means  of  oxyacetylene  and 
oxy hydrogen  torches. 

It  was  Priestley  who  first  thought  of  using  oxygen  to  produce  high 
temperatures.  He  found  that  blowing  pure  oxygen  on  a  piece  of 
glowing  wood  would  cause  it  to  burn  furiously.  A  few  years  later 
an  American  scientist,  Robert  Hare,  of  Philadelphia,  put  this  dis- 
covery of  his  friend  Priestley  to  practical  use  by  inventing  the  oxyhy- 
drogen  torch,  or  blowpipe. 

The  oxyhydrogen  torch  consists  of  two  tubes,  one  inside  the  other. 
Hydrogen  gas  passes  through  the  outer  tube  and  is  ignited  at  the  tip 
of  the  torch.  Pure  oxygen  passes  through  the  inner  tube  and  the 
mixture  of  the  two  gases  burns  at  the  tip  of  the  torch  with  an  ex- 
tremely hot  flame,  about  2400°C.,  a  temperature  much  higher  than 
the  melting  point  of  iron. 

The  oxyhydrogen  torch  was  never  widely  used.  Instead,  the  oxy- 
acetylene torch  is  used.  With  the  oxyacetylene  torch,  a  flame  tem- 
perature of  over  3300°C.  may  be  easily  produced.  The  oxyacetylene 
torch  is  similar  in  principle  to  the  oxyhydrogen  torch,  but  acetylene 

Fig.  4.  Blast  lamp.  Similar  in  principle 
to  the  oxyhydrogen  torch  except 
that  compressed  air  and  any  fuel 
gas  are  used  instead  of  oxygen  and 
hydrogen.  The  lamp  is  used  by  glass 
blowers  and  jewelers. 






gas  is  used  instead  of  hydrogen  (see  Fig.  3) .  The  oxygen  used  in 
the  torch  is  stored  under  high  pressure  in  strong  steel  cylinders. 
The  acetylene,  however,  is  not  under  high  pressure,  but  is  dissolved 
in  a  liquid  called  acetone.  In  almost  every  automobile  service  garage, 
oxyacetylene  torches,  with  their  accompanying  cylinders  ol  oxygen 
and  acetylene,  may  be  seen  ready  for  use. 

The  oxyacetylene  torch  is  an  important  industrial  tool.  It  is  used 
to  weld,  cut,  and  clean  metal.  It  is  also  used  in  heat  treating  sur- 
faces of  metal  machine  parts  to  make  them  more  wear  resistant. 

Oxygen  rusts  and  derusts  steel.  In  the  presence  of  air  and  minute 
quantities  of  water  vapor,  steel  rusts.  In  rusting,  oxygen  from  the 
air  unites  slowly  with  the  metal,  forming  a  brown,  scaly  oxide  of 
iron.  At  the  high  temperatures  used  in  the  making  of  steel,  rust  lorms 
very  rapidly  and  is  a  serious  problem.  As  red-hot  steel  is  carried  to 
the  rolling  mills,  it  becomes  covered  with  seams  of  iron  oxide. 

These  surface  imperfections  are  removed  by  a  process  called 
torch-deseaming,  or  scarfing.  In  the  scarfing  operation,  the  llame  of 
the  oxyacetylene  torch  is  directed  onto  the  hot  steel.  The  surface 
of  the  steel  is  quickly  oxidized  to  a  depth  of  about  one  quarter  of  an 
inch.  The  iron  oxide  falls  off  readily,  leaving  an  unblemished  sur- 
face. Scarfing  may  be  done  by  hand  or  by  special  machines. 

(left)  Red-hot  steel  slab  passes  through  scarfing  machine 
in  which  oxyacetylene  torches  remove  surface  defects, 
(right)  Welding  metal  plates  with  an  oxyacetylene  torch. 

Linde  Air  Products  Company 



Because  of  the  high  tem- 
perature of  the  torch,  an 
oxyacetylene  weld  is 
smooth  and  strong. 

/.nidc  Air  Product*  Company 

Oxygen  saves  lives.  Priestley  also  discovered  another  use  for  oxy- 
gen. Alter  inhaling  oxygen,  he  wrote:  "My  breath  felt  peculiarly 
light  and  easy.  It  (oxygen)  may  be  peculiarly  salutary  to  the  lungs  in 
certain  cases  where  the  common  air  is  not  sufficient." 

Today,  pure  oxygen  is  administered  to  persons  with  pneumonia 
and  in  other  cases  where  the  respiratory  system  cannot  function  at 
its  normal  rate.  Usually  about  two  gallons  of  oxygen  are  administered 
per  minute.  Since  air  contains  only  about  20  percent  pure  oxygen, 
a  patient  who  is  weak,  or  whose  lungs  are  congested  or  partially 
destroyed,  can  satisfy  the  oxygen  requirements  of  his  body  by  breath- 
ing a  much  smaller  volume  of  air  that  is  rich  in  oxygen  than  he 
would  normally  require  of  ordinary  air.  Oxygen  is  usually  admin- 
istered by  means  of  an  oxygen  tent,  a  canopy  which  fits  over  the 
patient's  bed.  Pure  oxygen  is  introduced  at  such  a  rate  that  the  air 
inside  the  canopy  always  contains  from  45  to  GO  percent  oxygen. 

Oxygen  in  determining  basal  metabolism.  Pure  oxygen  is  used  by 
physicians  in  determining  the  rate  at  which  a  person's  food  supply 
is  oxidized  while  the  person  is  at  rest.  This  is  known  as  his  rate  of 
basal  metabolism.  This  rate  is  obtained  by  measuring  the  volume  of 
oxygen  consumed  by  the  person  at  rest  during  a  short  interval, 
usually  eight  minutes. 

From  these  data  the  number  of  liters*  of  oxygen  consumed  per 
minute  may  be  calculated  easily.  This  number  is  then  compared 
with  the  basal  metabolic  rate  for  a  normal  person  of  the  same  age, 
sex,  height,  and  weight.  Persons  in  normal  health  use  oxygen  at  a 

*  A  liter  (I.)  is  a  unit  of  capacity  (volume)  in  the  Metric  System.  It  is 
slightly  larger  than  a  U.S.  liquid  quart. 



standard  rate.  In  certain  diseases,  the  patient's  rate  is  higher  and 
in  others  it  is  lower  than  known  standards  for  persons  in  good 
health.  For  example,  a  high  basal  metabolic  rate  always  accompanies 
an  overactive  thyroid  (see  pages  147,  581).  A  low  basal  metabolic 
rate  may  be  an  indication  of  an  underactive  thyroid. 

Oxygen  flies  high.  Aviators  and  mountain  climbers  who  ascend  to 
high  altitudes  where  the  atmosphere  is  very  thin  must  carry  supplies 
of  oxygen.  Otherwise  their  senses  become  dulled  and  they  are  likely 
to  lose  consciousness.  The  United  States  Air  Force  requires  the  use 
of  oxygen  at  altitudes  above  10,000  feet  in  the  daytime  and  from  the 
ground  up  at  nighttime.  Airliners  flying  at  altitudes  ranging  from 
15,000  to  more  than  .SO, 000  feet  contain  equipment  to  keep  the  oxy- 
gen concentration  inside  their  cabins  only  slightly  less  than  that  at  sea 
level.  The  cabins  are  airtight  and  are  pressurized  by  means  of  pumps, 
so  that  the  inside  pressure  does  not  become  uncomfortably  low. 

Rescue  parties  entering  mines  and  buildings  in  which  dangerous 
gases  are  present  carry  oxygen-breathing  apparatus. 

Air  Photographic  and  Charting  Service,  U.S.  Air  Force 

High  in  the  strato- 
sphere, this  pilot  is  de- 
pendent upon  a  con- 
tinuous supply  of  pure 


Some  other  uses  of  oxygen.  Oxygen  is  used  in  the  photoflash  lamps 
employed  in  photography.  These  lamps  look  like  ordinary  electric- 
light  bulbs,  but  they  are  filled  with  oxygen  and  aluminum  foil.  When 
an  electric  current  is  passed  through  the  filament,  the  aluminum  foil 
is  raised  to  its  kindling  temperature.  It  ignites  and  oxidizes  with 
a  blinding  flash.  This  Hash  lasts  only  about  ^  of  a  second. 

Pure  oxygen  has  also  been  introduced  in  the  manufacture  of  steel, 
synthetic  gasoline,  and  fuel  from  underground  coal  (see  pages  335, 
3(50,  417). 

How  is  such  a  huge  amount  of  oxygen  prepared?  It  is  not  sur- 
prising to  find  that  the  two  main  sources  of  oxygen  are  air,  the  most 
abundant  mixture  containing  oxygen,  and  water,  the  most  abundant 
compound  ol  oxygen.  Nearly  all  commercial  oxygen  is  obtained 
from  liquid  air  by  first  lowering  the  temperature  of  the  air  until  it  is 
changed  to  a  liquid.  The  preparation  of  oxygen  from  liquid  air  is 
discussed  on  page  100.  Less  than  one  percent  of  the  oxygen  produced 
commercially  is  obtained  by  decomposing  water  by  means  ol  an 
electric  current.  This  process,  called  the  electrolysis  of  water,  is  de- 
scribed on  page  62. 

How  can  we  test  for  oxygen?  When  Priestley  prepared  oxygen 
from  mercuric  oxide,  he  tested  the  gas  by  placing  burning  and 
glowing  substances  in  it.  In  each  case,  the  substance  burned  more 
vigorously.  This  method  is  still  used  to  identity  oxygen.  A  glowing 
splint  or  splinter  ol  wood  is  thrust  into  a  bottle  oi  the  gas.  Such  a 
splint  placed  in  a  bottle  of  oxygen  bursts  into  flame  at  once. 

No  other  odorless  gas  will  cause  a  glowing  splint  to  burst  into 
flame  in  this  way.  Hence,  we  can  distinguish  oxygen  from  any  other 
odorless  gas  by  this  simple  procedure.  We  call  such  a  method  of 
identifying  a  substance  a  test  for  that  substance.  Chemists  have 
devised  hundreds  of  tests,  which  they  use  in  identifying  many  other 
pure  substances. 

Ozone,  the  active.  Ten  years  after  Priestley's  discovery  of  oxygen, 
another  gas  which  possessed  a  peculiar  odor  and  which,  unlike 
oxygen,  tarnished  mercury  under  normal  conditions,  was  reported. 
But  it  was  not  until  1810  that  Schoenbcin  (shun'bfn)  isolated  this 
gas  and  called  it  ozone,  from  the  Greek  word  meaning  to  smell.  Its 
sharp  odor  is  noticeable  around  electric  machines  in  operation. 

O/one  is  a  pale  blue  gas,  one  and  one-half  times  as  heavy  as 
oxygen.  It  is  even  less  soluble  in  water  than  oxygen  but  is  more 
active  chemically.  It  is  a  strong  oxidizing  agent.  That  is,  it  is  a 
substance  which  readily  supplies  oxygen  for  chemical  union  with 
another  substance. 

/.  -' 

current  L'"11' 

Fig.  5.  A  continuous-process  ozone  tube.  Dry  air  enters 
at  lower  left,  passes  through  the  brush  discharge  of  high- 
voltage  current.  Part  of  the  oxygen  of  the  air  is  converted 
to  ozone  which  leaves  through  pipe  at  right. 

Ozone  (written  O8)  is  prepared  by  passing  electric  discharges 
through  cither  dry  air  or  oxygen  (written  O.,)  .  About  eight  percent 
of  the  oxygen  is  converted  into  pure  o/.one,  although  slightly  larger 
yields  are  obtained  if  the  temperature  is  kept  low,  or  if  the  process 

is  continuous. 

3  volumes  of  oxygen 

-  2  volumes  of  ozone 

Ozone  is  unstable;  it  changes  back  to  oxygen  quickly,  two  volumes 
of  ozone  changing  into  three  volumes  of  oxygen.  It  cannot  be  stored 
and  must  be  produced  at  its  point  of  use. 

What  is  allotropy?  As  we  have  just  learned,  oxygen  exists  in  two 
forms:  ordinary  oxygen  and  o/one.  The  existence  in  the  same 
physical  state  (both  oxygen  and  o/.one  are  gases)  of  two  or  more 
forms  of  the  same  element  is  a  phenomenon  called  allotropy 
(0-lot'r6-pi) .  The  various  allotropic  forms  of  an  element  have  differ- 
ent physical  and  chemical  properties. 

The  cause  of  allotropy  is  not  yet  completely  understood.  Hut  we 
know  that  it  is  caused  in  part  by  differences  in  the  arrangement  of 
the  atoms  and  in  the  amount  of  energy  in  the  various  allotropic 
forms.  This  can  be  seen  easily  by  referring  to  the  way  in  which  o/one 
is  produced  from  oxygen.  When  an  electric  current  discharges 
through  oxygen,  the  electric  energy  changes  oxygen  into  ozone, 
which,  as  you  would  expect,  possesses  more  energy  than  ordinary 
oxygen.  O/one,  on  changing  back  into  oxygen,  liberates  this  energy 
in  the  form  of  heat.  During  the  change  from  oxygen  to  ozone,  or 
from  o/one  to  oxygen,  no  energy  is  destroyed,  nor  is  any  energy  cre- 
ated. The  change  from  one  allotropic  form  of  oxygen  to  the  other 
is  an  excellent  example  of  the  law  of  the  conservation  of  energy. 



Allotropy  is  not  confined,  of  course,  to  oxygen.  For  example,  the 
element  phosphorus  occurs  in  two  allotropic  forms:  white  phos- 
phorus and  red  phosphorus.  White  phosphorus  melts  at  44°C.,  is 
poisonous,  soluble  in  carbon  disulfide,  and  has  a  very  low  kindling 
temperature.  Red  phosphorus  is  nonpoisonous,  insoluble  in  carbon 
disulfide,  and  has  a  higher  kindling  temperature  than  white  phos- 
phorus. Red  phosphorus  is  also  heavier  and  less  active  chemically 
than  white  phosphorus.  Under  proper  conditions,  red  phosphorus 
may  be  changed  into  white  phosphorus,  and  vice  versa. 

Ozone  "burns  up"  germs.  Because  of  its  extreme  chemical  activity, 
o/one  is  used  to  a  limited  extent  in  purifying  water.  It  kills  bacteria 
and  other  microorganisms  in  water  by  oxidi/ing  them,  literally  burn- 
ing them  up.  Other  organic:  materials  present  are  also  oxidized. 
About  one  gram*  of  ozone  will  purify  a  cubic  meter  of  water. 

Ozone  destroys  odors.  Because  o/one  is  an  excellent  oxidi/ing 
agent,  it  is  used  also  in  purifying  air  in  homes,  refrigerators,  tunnels, 
and  zoos.  Small  ultraviolet  lamps  change  the  oxygen  in  air  to  ozone 
which  clears  away  bad  odors.  Because  of  the  increasing  abundance  of 
low-cost  electricity,  it  seems  possible  that  o/one  may  be  used  more 
widely  in  the  removal  ot  unpleasant  tastes  and  odors  from  water 
than  it  now  is. 

Ozone  helps  screen  the  earth.  Ozone  is  present  in  the  layers  of 
the  atmosphere,  about  30  miles  above  the  surface  of  the  earth.  Sci- 
entists believe  that  this  region  which  is  rich  in  o/one  acts  as  a  screen 
that  protects  life  on  earth  from  the  harmful  effects  of  too  much  ultra- 
violet light  from  the  sun. 

What  part  does  chance  play  in  scientific  discoveries?  Some  years 
after  his  discovery  of  oxygen,  Priestley  commenting  on  this  memo- 
rable occasion  said:  "I  can  not  at  this  distance  of  time  recall  what  it 
was  that  I  had  in  view  in  making  this  experiment,  but  I  had  no 
expectation  of  the  real  issue  of  it.  If  I  had  not  happened  to  have  had 
a  lighted  candle  before  me,  I  should  probably  never  have  made  the 
trial,  and  the  whole  train  of  my  future  experiments  relating  to  this 
kind  of  air  might  have  been  prevented.  More  is  owing  to  what  we 
call  chance  than  to  any  proper  design  or  preconceived  theory." 

Chance  may  have  played  some  small  part  in  leading  Priestley  to 
make  his  experiments.  It  seems  likely,  though,  that  he  failed  to  take 
into  account  the  consuming  natural  curiosity,  always  present  in  true 
scientists,  which  literally  forced  him  to  make  his  experiments.  No 

*  A  gram  (g.)  is  a  small  unit  of  weight  or  mass  in  the  Metric  System.  One 
thousand  grams  are  equal  to  a  kilogram,  which  weighs  slightly  more  than  2.2 
pounds.  For  a  definition  of  mass,  see  page  157. 


doubt  one  of  the  tests  he  would  eventually  have  made  on  any  gas 
was  to  see  if  it  would  burn.  If  the  lighted  candle  had  not  been 
present  at  that  particular  moment,  he  undoubtedly  would  have 
lighted  one  later  on.  Probably  what  Priestley  meant  by  the  statement 
was  simply  that  he  had  no  specific  purpose  in  mind  in  making  the 
experiment.  But  having  no  specific  purpose  in  mind  and  attributing 
the  results  to  chance  are  not  the  same.  Nevertheless,  such  confusion 
exists  among  scientists  even  today. 

Other  men  were  performing  experiments  similar  to  those  made  by 
Priestley  (among  them  Scheele) ,  and  we  might  say  that  the  time  was 
ripe  for  these  discoveries.  If  Priestley  had  failed  to  make  them,  no 
doubt  some  other  experimenter  soon  would  have  made  them. 

Today  carefully  planned  research  has  speeded  up  advances  in 
science.  Chance  now  plays  a  very  slight  role  in  the  development  of 
chemistry.  Surely  we  cannot  regard  as  "chance"  the  keenness  of 
mind  to  appreciate  the  significance  ol,  and  to  follow  up  by  intelligent 
experiment,  a  clue  furnished  by  some  unforeseen  event. 

Progress  in  science  and  changes  in  society  are  closely  interrelated. 
Most  of  us,  although  we  live  in  a  world  ol  science,  have  strange 
notions  about  what  scientists  arc  like.  We  picture  the  research  man 
as  a  lone,  mysterious  genius  who  locks  himsell  up  in  his  laboratory 
away  from  the  world,  and  attempts  to  solve  some  abstract  scientific 
problem.  We  sec  him  emerging  triumphant,  after  weeks  or  perhaps 
years,  with  some  great  discovery.  Most  of  us  have  an  idea  that  the 
efforts  of  the  man  of  science  arc  influenced  very  little  by  the  society 
in  which  he  lives.  Nothing  could  be  farther  from  the  truth.  The 
scientist  is  influenced  by  society  and  society  in  turn  is  influenced 
by  the  scientist. 

The  steam  engine,  for  example,  was  developed  out  of  the  social 
needs  of  the  eighteenth  century.  With  the  Industrial  Revolution, 
which  began  in  Birmingham,  England,  came  a  need  for  more  iron  to 
make  the  machinery  so  much  in  demand.  In  making  iron,  wood  char- 
coal had  been  used,  but  England's  forest  reserves  had  been  seriously 
depleted  by  the  use  of  timber  in  the  many  ships  swallowed  up  in 
naval  wars  and  commercial  ventures.  As  a  result,  coal  had  to  be  sub- 
stituted for  charcoal,  and  coal  mines,  abandoned  because  they  were 
Hooded,  had  to  be  reopened.  This  meant  draining  mines,  and  James 
Watt  improved  Newcomen's  steam  engine  for  this  purpose. 

At  about  the  same  time,  the  problem  of  burning  was  reinvesti- 
gated,  and  Priestley's  investigations  (which  led  to  his  discovery  of 
oxygen)  were  closely  related  to  society's  need  for  more  information 
concerning  the  extraction  of  metals. 




Becker,  Carl.  Modern  History,  pp.  581-607.  Silver  Burdett 
Co.,  New  York,  1952.  "How  Science  Gave  Men  Machines  to 
Work  for  Them,  and  How  the  Machines  Changed  the  Con- 
ditions under  Which  Men  Had  to  Live  and  Labor."  Very 

Ficklen,  Joseph  B.  "Dust  Explosions."  Journal  of  Chemical 
Education, 'March,  1942,  pp.  131-134.  Published  by  the  Amer- 
ican Chemical  Society,  Easton,  Pa.  Editorial  office:  Metcalf 
Chemical  Laboratory,  Brown  University,  Providence,  R.I.  An 
interesting  and  well-illustrated  article. 

Friend,  J.  N.  Man  and  the  Chemical  Elements.  Scribners, 
New  York,  1953.  The  author  rambles  along  into  many  inter- 
esting sidepaths  and  offers  sidelights  such  as  the  origin  of 
gibberish  from  the  alchemist  Geber. 

Partington,  J.  R.  A  Short  History  of  Chemistry,  pp.  110-120. 
The  Macmillan  Co.,  London,  1939.  The  fascinating  story  of 
Priestley's  discovery  of  oxygen. 


1.  A  catalyst  is  a  substance  that  changes  the  speed  of  a 
chemical  reaction  without  being  itself  permanently  changed. 

2.  Oxidation  is   the  chemical   union  of  a  substance  with 
oxygen.  In  slow  oxidation,  neither  light  nor  noticeable  heat  is 
liberated.  In  rapid  oxidation,  light  and  noticeable  heat  are 
evolved.  Burning  is  rapid  oxidation. 

3.  The  kindling  temperature  ot  a  substance  is  the  lowest 
temperature  at  which   that  substance  catches   fire  and  con- 
tinues to  burn. 

4.  Spontaneous  combustion  is  a  burning  started  by  the  heat 
accumulated  during  slow  oxidation.  Combustion  is  any  chem- 
ical action  that  liberates  heat  and  light. 

5.  Finely  divided  powders  form  explosive  mixtures  with  air 
because  of  the  extremely  large  surfaces  exposed  to  oxygen. 

6.  Elements  may  occur  in  two  or  more  varieties,  or  allotropic 
forms,  in  the  same  physical  state.  These  allotropic  forms  differ 
in  their  properties  because  of  differences  in  the  arrangement 
of  their  atoms  and  differences  in  the  amount  of  energy  pos- 
sessed by  particles  of  each  allotrope. 

7.  Today  carefully  planned  research  has  speeded  up  ad- 
vances in  science,  and  chance  discoveries  play  a  smaller  part 
in  the  development  of  chemistry  than  they  once  did. 

8.  Scientific  progress  and  social  changes  are  interrelated. 



Group  A 

1.  Make  as  large  a  list  as  you  can  of  substances  containing 
oxygen  combined  with  other  elements. 

2.  From  what  chemical  compound  was  O2  first  prepared 
by  Priestley? 

3.  From  what  compounds  is  O2  usually  prepared  in  the 

4.  (a)  What  is  a  catalyst?   (b)  Illustrate  your  answer. 

5.  What  assurance  have  we  that  the  MnO.,  used  in  the  lab- 
oratory preparation  of  O2  acts  as  a  catalytic  agent? 

6.  What  happens  to  the  KCl  that  is  formed  by  the  de- 
composition of  KC1O3? 

7.  Why  is  it  possible  to  collect  O2  by  the  displacement  of 

8.  Priestley  for  a  time  collected  gases  by  the  displacement 
of  mercury.  Why  do  we  not  use  this  method  in  collecting  O.,? 

9.  List  five  physical  properties  of  O.,. 

10.  Discuss  the  most  important  chemical  property  of  O2. 

11.  (a)  What  is  slow  oxidation?  (b)  rapid  oxidation?  (c)  Il- 
lustrate each. 

12.  What  is  the  kindling  temperature  of  a  substance? 

13.  Devise  an  experiment  to  show  that  air  is  necessary  for 

14.  Name  four  substances  whose  kindling  temperatures  are 
lower  than  that  of  coal. 

15.  Why  is  air  removed  from  an  electric-light  bulb? 

16.  It  is  not  as  easy  to  burn  a  500-page  book  whole  as  it  is  to 
burn  the  same  book  a  page  at  a  time.  Explain. 

17.  Why  is  asbestos  used  in  making  theater  curtains? 

18.  (a)  What  are  the  conditions  necessary  for  spontaneous 
combustion?   (b)  Show  how  these  conditions  work  out  in  the 
spontaneous  combustion  of   (1)  finely  divided  coal  in  the  en- 
closed hold  oi  a  ship,  and   (2)  moist  hay  in  a  hayloft. 

19.  Why  may  the  presence  of  dust  in  the  air  of  a  Hour  mill 
cause  a  frightful  explosion? 

20.  Make  a  list  of  the  uses  of  Or 

21.  Distinguish  between  combustion  and  burning. 


22.  What  is  an  oxidizing  agent? 

23.  Describe  the  greatest  industrial  use  of  O2. 

24.  O2  rusts  steel.  What  part  does  O2  play  in  derusting  steel? 

25.  How  do  oxygen  tents  aid  patients  suffering  from  pneu- 
monia and  other  respiratory  diseases? 

26.  (a)  What  is  basal  metabolic  rate?    (b)  How  is  it  de- 
termined?   (c)  What  is  its  significance? 

27.  How  does  O2  play  an  important  part  in  aviation? 

.  .  T .  . . 

-  I  -  •  • 

28.  State  the  two  most  important  industrial  .sources  of  O0. 

29.  What  three  factors  must  be  considered  in  selecting  the 
method  for  preparing  large  quantities  of  a  substance? 

30.  (a)  Describe  a  test  tor  O2.  (b)  What  does  the  word  test 
mean  to  a  chemist? 

31.  O2  and  O3  can  be  changed  easily  from  one  to  the  other. 
(a)  How?   (b)  What  great  law  does  this  illustrate? 

32.  How  is  O3  prepared? 

33.  How  does  O8  differ  from  O2  in  physical  properties? 

34.  By  means  ot  an   illustration,  explain   the  meaning  of 
a  I  lot  ropy. 

35.  What  are  the  chief  uses  of  O3? 

36.  Write  word-equations  for  the  burning  of   (a)  coal   (car- 
bon) ,    (b)  iron,    (c)  phosphorus,  and    (d)  sultur. 

Group  B 

37.  By  using  mercury  Priestley  was  able  to  collect  a  number 
of  gases  that  had  escaped   the  attention  of  other  scientists. 

38.  Science  today  depends  less  upon  chance  discoveries  than 
it  has  in  the  past.  Explain. 

39.  Window  curtains  behind  a  fish  bowl  filled  with  water 
caught  fire.  Was  this  spontaneous  combustion?  Explain. 

40.  Why  is  O,  passed  through  the  inner  tube  ot  the  oxyhy- 
drogen  torch  rather  than  through  the  outer  tube? 

41.  Because  ot  oxidation,  linseed  oil  hardens  when  exposed 
to  air.  How  is  this  chemical  reaction  speeded  up? 

42.  Fishes  die  in  air,  man  drowns  in  the  sea.  Explain. 

43.  Blow  against  a  burning  candle  and  it  goes  out.  Blow  on 
the  slowly  dying  embers  of  a  fire  and  they  burn  more  actively. 

44.  Why  is  green,  or  moist,  hay  more  susceptible  to  spon- 
taneous combustion  than  dry  hay? 


45.  Illustrate  the  statement  "progress  in  science  and  changes 
in  society  are  closely  interrelated"  by  an  example  not  men- 
tioned in  this  chapter. 

46.  Coal  dust  is  sometimes  shoveled  into  a  burning  coal 
furnace.  Why? 


1.  Have  you  ever  heard  of  pure  oxygen  being  administered 
to  athletes  before  a  strenuous  game?  Consult  the  football  or 
track  coach  in  your  school  or  in  a  nearby  college.  What  do 
you  conclude?  Explain. 

2.  Organize  a  small  group  in  your  chemistry  class  and  with 
the  help  of  your  teacher  set  up  apparatus  such  as  Priestley  first 
used  when  he  obtained  oxygen.  A  large  burning  glass  can  be 
borrowed  from  the  physics  department.  Demonstrate  the  ex- 
periment before  your  class. 

3.  Watch  your  mother  put  up  fruit  and  jam  in  jars.  De- 
scribe her  procedure,  and  explain  why  the  food  is  heated, 
and  why  a  layer  of  paraffin  is  placed  over  the  food  before  the 
jar  is  sealed. 

.4.  Consult  your  doctor  and  a  good  book  on  first  aid,  and 
write  a  report  on  the  diagnosis  and  treatment  of  first-,  second-, 
and  third-degree  burns. 

5.  Read  the  story  of  a  recent  scientific  discovery,  (a)  From 
the  account  you  read,  what  part  do  you  conclude  that  chance 
plays  in  scientific  discoveries?  (b)  Why  may  the  scientist 
himself  believe  chance  plays  a  more  important  part  than  it 
actually  does? 

H  Y  D  R  O  G  E 


.  .  .  Cavendish  was  almost  passion- 
less. An  intellectual  head  thinking, 
a  pair  of  wonderful  acute  eyes  ob- 
serving, and  a  pair  of  skilful  hands 
*  •.  experimenting  or  recording  are  all 

that  I  realize  in  reading  his  memoirs. 
Cavendish  did  not  stand  aloof  from 
other  men  in  proud  spirit,  he  did 

>  ,  ,  so  conscious  of  his  inferiority,  not 

boasting  of  his  excellence. 

Dr.  George  Wilson,  1851 

An  eccentric  man  of  science.  While  Priestley  was  performing  his 
immortal  experiments,  another  Englishman  was  puttering  around 
in  his  palace  laboratory  on  the  wandering  trail  of  phlogiston.  This 
man  was  Henry  Cavendish,  one  of  the  most  eccentric  persons  in  the 
whole  history  of  science.  It  was  said  that  he  was  "the  richest  among 
the  learned  and  the  most  learned  among  the  rich."  Although  the 
richest  man  in  all  England,  he  shut  himself  up  in  his  private  labora- 
tory and  spent  more  than  60  years  tracking  down  many  of  the  secrets 
of  nature. 

In  addition  to  his  discoveries  in  chemistry,  Cavendish  was  inter- 
ested in  physical  problems.  His  work  in  the  fields  of  heat  and  elec- 
tricity was  of  highest  rank;  later  it  was  followed  up  by  the  work  of 
other  eminent  scientists  such  as  Joseph  Black  and  Michael  Faraday. 
The  elusive  trail  of  phlogiston.  Two  hundred  and  fifty  years 
before  the  work  of  Cavendish,  Paracelsus  (par-d-sel'sus)  of  Switzer- 
land had  noticed  bubbles  of  gas  rise  from  an  acid  into  which  iron  is 
dropped.  Although  he  found  that  this  gas  burns,  he  carried  his 
investigations  no  further.  Then  came  Cavendish  to  whom  the  search 
for  truth  was  the  ruling  motive  of  life.  He  too  noticed  the  gas 
evolved  when  zinc  or  iron  is  dropped  into  an  acid  and  went  to  work 



to  investigate  this  phenomenon.  He  collected  the  gas  carefully,  and 
made  a  thorough  study  of  it.  He  named  it  inflammable  air  because  it 
burned.  He  thought  he  had  obtained  phlogiston  itself. 

Strangely  enough,  the  discovery  of  this  gas,  coupled  with  the  dis- 
covery of  oxygen,  paved  the  way  for  the  complete  overthrow  of  the 
phlogiston  theory  and  the  establishment  ol  l.axoisier's  true  explana- 
tion of  burning.  Though  Priestley  died  still  believing  in  the  phlo- 
giston theory,  Cavendish,  when  the  discussion  over  Lavoisier's  new 
chemistry  became  very  heated,  gave  up  his  active  interest  in  chemistry 
and  turned  to  the  problem  of  determining  the  weight  of  the  earth. 
He  said  that  he  had  no  patience  with  squabbles,  that  he  was  inter- 
ested in  experimentation,  not  controversy. 

How  is  hydrogen  produced  in  the  laboratory?  Cavendish's  lab- 
oratory method  of  preparing  hydrogen  is  still  used.  Zinc  is  placed 
in  a  generator,  as  shown  in  Fig.  6,  and  dilute  hydrochloric  or  sul- 
turic  acid  is  poured  over  it  through  a  thistle  tube.  Bubbles  of  hydro- 
gen gas  form  at  once,  and  heat  is  generated.  The  hydrogen  is  col- 
lected in  the  same  way  as  oxygen,  that  is,  by  the  displacement  of 
water.  Why  is  hydrogen  not  collected  by  the  displacement  of  air? 

The  chemical  change  that  takes  place  is  represented  as  follows: 

Zinc  -4-  hydrochloric  acid  • 
Zn   +  2HC1 

>  zinc  chloride  4-  hydrogen 
ZnCl2         4-        H2 

Hydrochloric  acid  is  a  compound  of  hydrogen  and  chlorine.  The 
zinc  takes  the  place  of,  or  replaces,  the  hydrogen  of  the  acid  and 
liberates  it  as  a  free  gas.  Instead  of  hydrochloric  acid,  we  now  have 
/inc  chloride,  which  remains  dissolved  in  the  water  in  the  generator. 
Zinc  chloride  is  a  white  solid. 

All  acids  contain  hydrogen,  which  may  be  replaced  by  certain 
metals.  Hence,  to  prepare  hydrogen,  we  may  use  almost  any  acid  and 
one  of  a  number  of  other  metals  instead  of  zinc.  It  is  a  curious  fact 
that  when  pure  zinc  is  added  to  pure  hydrochloric  acid,  the  chemical 

thistle  tube 


Zn  +  HCI 


Fig.  6.  Laboratory  prepa- 
ration of  hydrogen.  Why 
does  thistle  tube  extend 
below  surface  of  liquid 
in  the  generator? 

hydrogen  gas 
collecting  jar 

-  trough  of  water 




Fig.  7.  Preparation  of  hydrogen 
by  the  action  of  sodium  on  water. 
Perform  carefully.  The  reaction 
may  be  violent. 

action  is  very  slow.  However,  slightly  impure  zinc  replaces  the 
hydrogen  of  the  acid  rapidly.  The  impurity  in  the  zinc  acts  as  a 

Can  hydrogen  also  be  prepared  from  water?  Perhaps  you  have 
been  thinking,  "II  the  hydrogen  of  an  acid  can  be  replaced  by  a 
metal,  can  the  hydrogen  of  water  also  be  replaced  by  a  metal?"  The 
answer  is  Yes.  Very  active  metals,  such  as  sodium,  potassium,  and 
calcium,  possess  enough  chemical  energy  to  replace  the  hydrogen  of 
water.  This  experiment  may  be  tried  by  filling  a  bottle  with  water 
and  inverting  it  in  a  dish  partly  filled  with  water,  as  shown  in  the 
illustration  (Fig.  7)  .  Wrap  in  filter  paper  a  piece  of  sodium  the 
si/.e  of  a  pea  and  put  the  paper  in  the  coiled  end  of  a  wire.  Then 
quickly  insert  this  end  of  the  wire  into  the  bottle.  Immediately,  a 
very  active  evolution  of  gas  is  noticed  and  the  bottle  becomes  filled 
with  a  colorless  gas.  The  equation  tor  this  reaction  is: 

Sodium  -f  water  — » sodium  hydroxide  -f  hydrogen 
2Na     +2HOH->  2NaOH  +         H2 

(Water  may  be  represented  as  H2O  or  HOH.) 

In  the  chemical  reaction  that  occurs,  sodium  replaces  half  of  the 
hydrogen  in  water,  and  forms  sodium  hydroxide,  NaOH.  (The  OH 
group  is  called  the  hydroxyl  group.)  The  sodium  hydroxide  remains 
dissolved  in  the  water.  The  water  solution  of  sodium  hydroxide  feels 
soapy,  and  turns  pink  litmus,  a  vegetable  coloring  matter,  blue.  By 
evaporating  the  solution,  the  sodium  hydroxide  can  be  separated 
from  the  water  as  a  white  solid.  Sodium  hydroxide  belongs  to  a  group 
of  compounds  called  bases  (Chapter  14) . 

Even  iron,  which  is  not  nearly  as  active  as  sodium,  will  replace  the 
hydrogen  of  water  while  the  water  (as  steam)  is  passed  over  red-hot 



iron  in  a  heated  tube.  In  tact,  this  method  ol  prcp.ning  hvdrogen 
has  been  used  to  a  slight  extent,  the  products  ol  ihe  ic.ution  being 
iron  oxide  and  hydrogen. 

What  is  the  industrial  method  ol  preparing  hydrogen?  In  the 
order  ol  the  quantities  ot  gas  produced,  hydrogen  is  obtained  lot 
commercial  use:  (1)  by  passing  steam  oxer  hot  carbon,  (2)  by  pass 
ing  steam  through  natural  gas  (methane)  ,  in  the  presence  ol  a 
catalyst,  (3)  by  the  electrolysis  ol  \\ater.  The  equations  tor  these 
methods  are: 

1)  Steam  -f  carbon  — >  carbon  monoxide  -f  hydrogen 

H«'9 ! !+."".".".  ?.:    -»  CO  +        Hj 

2)  Steam  -f  methane  — »  carbon  monoxide  -f  hydrogen 

H>O    +      CH4      -*  CO  +       3H, 

3 "i   Water  — >  hydrogen  -f  oxygen 
2H»O->      2H2        4-      O> 

In  the  first  two  methods,  the  hydrogen  mav  be  sepaiated  from 
tlie  carbon  monoxide  gas  by  chilling  the  mixtuie  ol  gases.  1  lie  cai- 
bon  monoxide,  whose  free/ing  point  is  much  higher  than  that  of 
hydrogen,  solidifies,  leaving  nearly  pure  hydrogen  gas  behind.  Only 
a  minor  part  ol  the  total  hydrogen  production  is  by  electrolysis. 

Physical  properties  of  hydrogen.  Hydrogen  resembles  oxygen  in 
most  ol  its  physical  characteristics.  It  is  a  colorless,  odorless  gas, 
slightly  soluble  in  water,  and  very  difhcult  to  liquely.  since  it  changes 
Irom  a  gas  to  a  liquid  at  — 252°C.  It  differs  physically  from  oxygen 
chiefly  in  its  weight.  It  is  the  lightest  element  known.  It  is  y1^  as 
heavy  as  oxygen  and  ^^  as  heavy  as  air.  One  liter  ot  hydrogen 
weighs  approximately  0.09  gram  at  standard  conditions' of  tempera- 
ture and  pressure. 

The    metals   palladium   and   platinum   absoib   large   volumes   ol 

hydrogen  gas 
porous  cup 

Fig.  8.  Passage  (diffusion)  of 
hydrogen  through  a  porous 
cup.  Why  is  water  forced 
through  the  glass  tube? 

glass  tube 



hydrogen.  This  absorption,  or  occlusion,  is  accompanied  by  such  an 
increase  in  temperature  that  the  metals  actually  glow.  The  absorp- 
tion of  hydrogen  by  these  two  metals  takes  place  in  one  type  of 
lighting  apparatus  sometimes  used  in  lighting  the  burners  of  gas 
stoves.  Such  a  lighter  usually  consists  of  a  fine  wire  or  wires  of  one 
of  the  two  metals  strung  between  two  suspension  points.  The  lighter 
is  placed  above  the  burner.  When  the  gas  is  turned  on,  hydrogen  is 
absorbed  by  the  wire.  This  causes  the  wire  to  give  off  heat.  In  about 
a  second,  the  kindling  temperature  of  the  gas  is  reached,  and  the  gas 
is  ignited. 

If  two  or  more  gases  are  at  the  same  temperature,  the  particles  of 
the  lighter  gas  move  more  rapidly  than  the  particles  of  the  heavier 
gases.  Since  hydrogen  is  the  lightest  gas  known,  its  particles  are  in 
very  rapid  motion,  passing  through  porous  substances  rapidly,  as 
shown  in  Fig.  8. 

Hydrogen  may  burn  quietly  or  explode  violently.  Although  pure 
hydrogen  burns  quietly  in  air  or  in  oxygen  with  a  pale  blue,  almost 
colorless,  flame,  a  mixture  of  hydrogen  and  oxygen  may  unite  with 
explosive  violence.  The  two  gases,  when  mixed  and  kept  below  a 
temperature  of  about  800°C.,  will  not  unite;  but  a  spark,  a  flame,  or 
a  temperature  of  above  800°C.  will  cause  them  to  unite  violently.  For 
this  reason,  great  care  must  be  taken  while  experimenting  with 
hydrogen  to  keep  all  flames  away  from  the  generator.  It  is  also  nec- 
essary to  wait  until  the  air  has  been  completely  expelled  from  the 
generator^  In  -lore  setting  fire  to  the  hydrogen  as  it  escapes  from  the 
delivery  wibe.  , 

When  the  Frenchman,  Pilatre  de  Rozier,  heard  of  this  gas  which 
Cavendish  had  studied,  he  tried  an  unusual  and  foolish  experiment. 

Fig.  9.  Oxidation  of  hydrogen  to  form  water.  Will  water  form 
if  bell  jar  becomes  hot?  How  may  bell  jar  be  kept  cool? 



He  inhaled  the  gas  until  he  had  filled  his  lungs,  and  then  as  the 
gas  issued  from  his  mouth  he  set  (ire  to  it.  All  Paris  held  its  sides 
with  laughter  as  it  watched  him  spitting  lire.  However,  when  he  set 
fire  in  the  same  way  to  a  mixture  of  this  gas  and  air.  "the  conse- 
quence was  an  explosion  so  dreadful  that  he  imagined  his  teeth  were 
all  blown  out." 

The  chemical  union  of  hydrogen  and  oxygen  may  be  written: 

2H2  -I-  O2  ->  2H2O 

This  is  an  example  of  the  strange  behavior  of  chemical  elements. 
Hydrogen,  a  highly  flammable  gas,  unites  with  oxygen,  a  gas  which 
helps  things  burn,  forming  water,  a  liquid  that  is  one  of  the  greatest 
enemies  of  fire. 

Experimental  proof  that  water  forms.  Of  course,  when  hydrogen 
combines  with  oxygen,  we  do  not  see  the  formation  of  a  flood  of 
water.  We  do  not  because  the  water  that  is  formed  at  the  tempera- 
ture of  burning  hydrogen  is  invisible  since  it  is  in  the  form  of  water 
vapor.  However,  it  is  possible  to  show  the  actual  formation  of  water 
by  arranging  the  apparatus  shown  in  the  illustration  (Fig.  9) .  As 
the  invisible  water  vapor  strikes  the  cool  surface  of  the  jar,  drops  of 
a  liquid  form.  This  liquid  is  pure  synthetic  water.  A  synthetic  com- 
pound is  a  compound  built  up  from  simpler  substances. 

Hydrogen  is  a  powerful  reducing  agent.  Since  hydrogen  has 
such  a  strong  attraction,  or  affinity,  for  oxygen,  it  is  able  to  tear 
oxygen  away  from  the  other  elements  of  many  of  its  compounds. 
The  removal  of  oxygen  from  a  compound  is  a  process 
known  as  reduction.  This  ability  of  hydrogen  to  remove  pxygen 
from  the  other  elements  of  a  compound  makes  it  a  reducing  or 

deoxidizing  agent. 

Fig.  10.  Reduction  of  copper  oxide  by  hydrogen. 

The  fishtail  burner  is  used  to  spread  the  flame, 
thus  heating  a  larger  surface  of  copper  oxide. 
What  is  the  function  of  the  drying  tube?  Why  is 
water  a  product? 

hydrogen  generator 




For  example,  if  pure  hydrogen  is  passed  over  black  copper  oxide 
brought  to  a  red  heat,  as  shown  in  Fig.  10,  the  hydrogen  takes  the 
oxygen  away  from  the  copper  oxide,  leaving  copper.  This  change 
may  be  represented  as  follows: 

Copper  oxide  (black)  +  hydrogen 
CuO  4-         H2 

>  water  4-  copper  (red) 
»  H2O     -f  Cu 

Although  hydrogen  is  a  good  reducing  agent  and  is  used  to  reduce 
the  oxides  of  such  metals  as  wolfram  and  molybdenum,  other  re- 
ducing agents  are  of  greater  commercial  use.  Perhaps  the  best 
example  of  these  is  carbon  (coke)  ,  which  is  used  in  reducing  iron 
ore  (iron  oxide)  to  iron. 

The  relation  of  oxidation  to  reduction.  In  the  experiment  just1 
described,  the  copper  oxide,  CuO,  is  reduced  to  copper.  At  the  same 
time  hydrogen  is  oxidi/ed  to  water,  H2O.  This  illustrates  a  general 
principle:  namely,  that  the  reducing  agent  is  itself  always  oxidi/ed. 
You  should  remember  that  whenever  one  substance  is  reduced, 
another  is  oxidi/.ed.  Thus  oxidation  and  reduction  always  occur  in 
the  same  reaction. 

The  bunsen  burner  and  how  it  works.  One  of  the  most  familiar 
pieces  of  apparatus  in  the  chemical  laboratory  is  the  bunsen  burner 
(named  after  Robert  Bunsen,  a  German  chemist  who  introduced  it  in 
1855)  .  The  function  of  this  burner  is  to  mix  a  gaseous  fuel  with 
air  in  order  to  make  a  lumliiminous  flame  that  has  a  high  tempera- 
ture and  will  not  deposit  soot. 

The  parts  of  the  bunsen  burner  designed  for  burning  manufac- 
tured gas  are  shown  in  Fig!  1  1.  Gas  enters  through  the  side  of  the 
stand,  and  its  speed  is  increased  by  passing  through  the  narrow  spud. 
The  rapidly  moving  gas  draws  in  air  through  the  hole  in  the  collar, 
which  may  be  adjusted  to  permit  the  correct  volume  of  air  to  enter. 
The  mixture  of  gas  and  air  passes  up  through  the  barrel  and  is 
ignited  at  the  top. 

When  the  pressure  of  the  gas  is  low,  the  flame  may  be  drawn  back 
and  may  burn  at  the  spud.  This  striking  back  of  the  flame  can  be 

Fig.  II. 

of     m 








•  It* 



dangerous  because  the  stand  -of  the  burner  becomes  hot.  enough  to 
produce  serious  burns  when  touched,  and  occasionally  the  hot  burner 
melts  the  rubber  hose  and  ignites  the  gas. 

In  the  type  of  burner  for  natural  gas,  the  air  intake  is  larger 
and  the  hole  in  the  spud  is  much  smaller.  It  is  also  supplied  with 
a  flame  retainer  which  has  0  small  pilot  jets  led  by  gas  bled  from 
the  main  tube.  This  eliminates  the  tendency  tor  the  flame  to  blow 

The  burner  on  a  gas  range  consists  of  a  series  of  small  bunsen 
burners.  Not  only  is  the  nonluminous  ilamc  of  the  bunsen  burner 
used  in  cooking,  but  it  is  used  also  to  some  extent  in  illumination  by 
burning  the  gas  inside  gas  mantles  consisting  of  09  pen  cut  thorium 
oxide  and  1  percent  cerium  oxide.  These  oxides,  when  heated,  glow 
with  a  rich  white  flame.  They  were  first  made  commercially  practical 
by  von  Welsbach  (ton  vels'baiO  .  The  widespread  use  of  electricity 
as  an  illimiinam  has  made  this  type  of  mantle  burner  all  but  obso- 

The  nature  of  a  flame.  A  flame  is  produced  only  when  combustible 
vapors  reach  their  kindling  temperature  in  the  presence  of  air, 
oxygen,  or  some  other  substance  that  supports  combustion.  When 
iron  is  heated  gently,  it  glows,  but  a  flame  is  not  produced,  because 
iron  does  not  vapori/e  at  low  temperatures.  A  candle,  on  the  other 
hand,  burns  with  a  flame,  because  the  candle  melts  and  the  heat 
from  the  burning  wick  causes  this  licjuid  to  change  to  a  vapor  that 
burns  in  air.  It  can  be  shown  that  the  interior  of  a  candle  flame 
consists  of  a  combustible  gas  by  holding  one  end  of  a  piping  hot 
tube  in  the  innermost  part  of  the  flame  and  then  touching  the  other 
end  with  a  lighted  match  (Fig.  12) . 

The  structure  of  the  bunsen  flame.  Three  distinct  /ones  are  no- 
ticeable in  the  bunsen  burner  flame  as  well  as  in  the  candle  flame. 
The  innermost  zone  A  is  composed  of  combustible  gas  that  has  not 
yet  reached  its  kindling  temperature.  This  fact  may  be  tested  by 
placing  a  match  head  in  zone  //.  This  may  be  clone  by  piercing 
a  match  near  the  head  with  a  pin.  The  pin  serves  as  a  bridge  across 

burning  gas 

Fig.  12.  Experiment  thawing 
j  the  nature  of  the  innermost 
:  zone  of  a  candle  flame. 



the  opening  of  the  burner.  If  the  burner  is  lighted  carefully,  the 
matcli  will  not  catch  fire.  Outside  zone  A  is  zone  B  in  which  the  gas 
is  burning.  The  light-purple  zone  C  is  the  region  of  complete  com- 
bustion, where  carbon  dioxide  and  water  vapor  are  formed. 

Just  below  the  tip  of  the  flame,  plenty  of  air  is  available  and, 
therefore,  this  is  the  oxidizing  part  of  the  flame.  Zone  B  is  the  region 
of  somewhat  incomplete  combustion,  where  one  of  the  products 
formed  is  carbon  monoxide,  a  reducing  agent.  Hence  this  zone  is 
the  reducing  part  of  the  flame.  This  zone  of  the  flame  is  used  in 
reducing  metallic  oxides  to  free  metals  by  means  of  carbon.  The 
blowpipe  directs  the  flame,  as  shown  in  Fig.  71,  page  321. 

The  luminosity  of  the  flame  of  the  bunsen  burner,  when  the  collar 
is  closed,  is  caused  by  the  decomposition  by  heat  of  a  small  amount 
of  the  hydrocarbons  and  the  subsequent  burning  of  free  carbon.  This 
may  be  proved  by  holding  a  cold  dish  in  the  luminous  flame.  The 
carbon  is  reduced  below  its  kindling  temperature  by  contact  with  the 
dish  and,  hence,  deposits  soot  on  the  dish. 

Hydrogen  rides  the  winds.  Interesting  historically,  but  of  only 
slight  commercial  importance  today,  is  the  use  of  hydrogen  in  filling 
balloons  and  other  lighter-than-air  craft.  Soon  after  the  discovery 
of  hydrogen.  Dr.  Charles  (shiirl) ,  of  Paris,  constructed  the  first 
large  hydrogen-filled  balloon,  and  in  the  presence  of  300,000  spec- 
tators, Pilatre  de  Ro/ier,  who  had  experimented  unwisely  with 
hydrogen  before,  bravely  climbed  inside  its  basket  and  started  on 
the  first  aerial  voyage  ever  made  by  a  human  being.  Since  that  time, 
balloons  filled  with  hydrogen  have  carried  men  around  the  world 
and  have  lifted  explorers  of  the  atmosphere  to  altitudes  of  almost 
14  miles.  The  development  of  the  motor-propelled,  rigid  balloon 
called  a  dirigible  gave  us  gigantic  craft,  nearly  1000  feet  long  and 
weighing  more  than  100  tons. 

From  "Photographic  History  of  the  Civil  War"  by  Albert  Shaw 

The  observation  balloon 
was  used  to  advantage  by 
the  Union  Army  in  the  War 
Between  the  States.  Here 
the  balloon  Intrepid  is  being 
filled  with  hydrogen  by  the 
generators  at  the  left. 

Such      helium-filled      bal- 
loons, carrying   recording! 
•  instruments,  are    used  to 
determine  weather  condi- 
tions at  high  altitudes. 

Official  Unttcd  States  A'nrj/  J'/mf,, r ,-,,,,-, 

The  history  of  this  type  of  flying  machine  was  filled  with  tragedy, 
to  a  large  extent  because  hydrogen  ignites  and  explodes  readily. 
With  the  discovery  of  large  sources  of  a  nonflammable  gas,  helium, 
and  its  subsequent  use,  this  ever-present  danger  was  removed.  There 
are  no  dirigibles  in  existence  today,  although  blimps,  which  are 
motor-propelled,  nonrigid  balloons,  are  used  for  military  purposes. 

Helium,  which  possesses  about  93  percent  of  the  lilting  power  of 
hydrogen,  may  be  obtained  today  at  a  cost  not  much  greater  than 
that  of  hydrogen.  The  United  States  is  fortunate  in  possessing  the 
largest  supply  of  helium  of  any  nation.  As  high  as  seven  percent 
helium  is  extracted  from  the  natural  gas  of  some  western  fields. 

To  conserve  our  own  supplies  and  prevent  other  nations  from 
using  helium  for  military  purposes,  Congress  passed  a  law  in  1938 
which  placed  its  export  under  very  strict  government  control.  Sci- 
entific discoveries  often  turn  out  to  be  two-edged  swords,  and  the 
careful  control  of  such  discoveries  should  be  part  of  the  business 
of  government. 

Hydrogen,  a  gas,  helps  solidify  fats  and  oils.  One  of  the  most 
important  uses  of  hydrogen  is  based  on  the  discovery  that  hydrogen 

Procter  and  Gamble  Compan\ 

These  machines,  called 
"freezers,"  are  used  in  the 
hydrogenation  of  vege- 
table oils  to  give  the  oils 
creaminess  and  smooth 



can  be  chemically  combined  with  other  substances  to  form  new 
products  of  great  value.  This  process  of  chemically  combining  hydro- 
gen with  other  substances  is  called  hydrogenation.  In  normal  times, 
hydrogenaiion  ol  fats  and  oils  is  the  largest  use  of  hydrogen. 

Generally,  during  hydrogenation,  liquid  oils  are  changed  to  semi- 
solid  fats.  Thus,  cottonseed  oil,  in  the  presence  of  a  finely  divided 
nickel  catalyst,  unites  with  hydrogen,  forming  a  white  fat  that  is 
often  sold  nuclei  such  trademarks  as  "Crisco"  or  "Spry."  Fish  oils, 
formerly  almost  useless,  have  been  hydrogenated  similarly  and  ren- 
dered useful.  Fats  for  soap-making  and  candle-making  have  been 
prepared  commercially  by  this  process  also.  Crude  oil,  coal,  and  cer- 
tain waste  products  of  petroleum  refining  are  also  hydrogenated  and 
changed  in  this  way  into  high-quality  gasoline.  Wood  alcohol  is  an- 
other important  chemical  made  with  the  aid  of  hydrogen. 

Oleomargarine,  or  margarine,  a  widely  used  butter  substitute,  is 
made  of  either  hydrogenated  vegetable  oils  or  animal  fats,  or  mix- 
tures of  these  substances.  Much  of  the  margarine  now  marketed  is  of 
vegetable  origin. 

The  introduction  of  hydrogenation  on  a  large  scale  helped  the 
cotton  farmers  of  the  South  by  opening  a  new  market  for  their  cot- 
tonseed oil.  It  has  had  even  greater  effects  abroad,  for  it  has  enabled 
several  Kuropean  countries  to  become  less  dependent  upon  other 
nations  for  their  vital  supplies  of  gasoline  and  edible  fats. 

Hydrogen  for  the  synthesis  of  ammonia.  Ammonia,  a  compound 
of  nitrogen  and  hydrogen,  leads  a  double  life.  It  is  an  important 
component  of  nearly  all  of  the  world's  most  powerful  explosives, 
and  it  is  an  important  ingredient  in  fertili/ers  used  all  over  the 
world.  The  most  significant  industrial  process  for  the  preparation  of 
ammonia  is  one  that  embodies  the  direct  combination,  or  synthesis, 
of  nitrogen  and  hydrogen.  This  process  is  discussed  in  detail  in 
Chapter  17.  In  wartime,  production  of  ammonia  probably  requires 
more  hydrogen  than  any  other  industrial  use. 

In  this  plant,  bituminous  coal  it 
hydrogenated  to  produce  many 
useful  chemicals. 

An  illustration  of  a  solar  prominence.  Solar  prominences  are  great  tongues  of 
glowing  hydrogen  which  shoot  out  of  the  chromosphere  of  the  sun  and  extend 
far  into  space.  The  flames  often  attain  lengths  of  more  than  100,000  miles  and 
have  been  known  to  reach  lengths  of  more  than  1,000,000  miles. 

Hydrogen  in  heating  and  cooling.  Hydrogen  has  another  impor- 
tant use.  When  it  is  burned,  it  gives  off  four  times  as  much  heat  as 
an  equal  weight  of  coal.  Because  of  the  high  temperature  produced, 
hydrogen  is  used  both  in  the  oxy hydrogen  torch  and  as  a  constituent 
of  certain  gaseous  fuels  such  as  water  grw,  which  contains  about  50 
percent  hydrogen.  Hydrogen  gas  is  also  used  as  a  cooling  agent  in- 
stead of  air.  Because  of  its  very  low  density  it  cuts  down  friction 
in  machines  such  as  high  speed  turbine-generators.  When  hydrogen 
is  so  used,  the  machine  is  completely  enclosed  to  shut  out  air. 

The  test  for  hydrogen.  Hydrogen  is  not  the  only  colorless  gas  that 
burns  with  a  pale  blue,  almost  invisible  flame.  But  it  is  the  only  gas 
that  forms  water  as  the  only  product  of  its  burning.  This  fact  gives 
us  a  simple  chemical  test  for  hydrogen. 

Where  is  hydrogen  found?  Unlike  oxygen,  only  small  amounts  of 
hydrogen  occur  on  earth  in  the  free  state.  However,  live  hydrogen 
is  the  most  abundant  element  in  the  sun.  The  immense  luminous 
tongues,  or  solar  prominences,  some  of  which  extend  half  a  million 
miles  from  the  sun's  surface,  consist  of  glowing  hydrogen.  It  is  also 
the  commonest  element  found  in  interstellar  space,  and  by  far  the 
most  abundant  material  out  of  which  the  whole  universe  is  built. 

Combined  hydrogen  is  very  common  on  earth.  Hydrogen  consti- 
tutes about  11  percent  by  weight  of  all  water  and  is  one  of  the  ele- 
ments in  petroleum,  all  acids,  and  living  cells  (protoplasm).  In 
spite  of  its  widespread  occurrence,  the  extreme  lightness  of  hydrogen 
accounts  for  the  fact  that  it  constitutes  only  1  percent  by  weight  of 
the  earth. 


Henry  Cavendish  (1731-1810),  one  of 
the  most  unusual  personalities  in  the 
history  of  science.  For  the  most  part,  he 
shunned  society  of  any  kind,  even  in- 
structing his  servants  to  keep  out  of  his 
sight.  Yet,  over  his  long  lifetime,  he 
made  many  important  contributions  in 
the  fields  of  chemistry  and  physics. 

Theories  lead  to  great  discoveries.  In  1932,  three  Americans 
headed  by  Harold  C.  Urey  discovered  that  ordinary  hydrogen  could 
be  separated  into  two  distinct  forms.  They  named  the  heavier  of 
these  forms  deuterium.  Two  years  later,  it  was  proved  that  there 
is  a  third  form  of  hydrogen.  This  form,  the  heaviest  of  the  three, 
has  been  named  tritium.  The  three  forms  have  the  same  chemical 
properties,  but  differ  in  certain  physical  properties. 

This  achievement  is  remarkable,  not  so  much  because  ordinary 
hydrogen  was  shown  to  be  a  mixture  of  three  forms  of  the  same 
element,  but  because  it  afforded  a  definite  example  of  a  great  con- 
tribution to  chemistry  made  by  scientists  who  forecast  these  dis- 
coveries purely  from  theory.  Their  success  points  to  the  fact  that 
science  needs  the  man  who  experiments,  the  thinker  who  can  put 
his  imagination  and  reason  to  work  in  propounding  theories,  and  the 
engineer  who  works  to  discover  how  these  theories  and  processes 
may  be  used  in  the  service  of  man. 

Another  example  of  how  theoretical  problems  in  science  may  turn 
out  to  be  of  great  practical  value,  was  the  theorizing  of  the  great 
mathematician,  James  Clerk  Maxwell.  In  1863  he  came  to  the  con- 
clusion that  just  as  light  results  from  a  wave  disturbance  in  the 
ether,  so  electric  disturbances  from  a  spark  should  produce  similar 
waves,  invisible,  to  be  sure,  but  nevertheless  existent.  Experimental 
evidence  of  such  waves  was  found  23  years  later  by  a  young  physicist, 
Heinrich  Hert/.  These  Hertzian  waves,  now  known  as  radio  waves, 
were  later  used  by  Marconi  in  the  transmission  of  wireless  messages. 
Thus  modern  radio  originated,  and  out  of  a  purely  theoretical  in- 
vestigation came  one  of  the  most  practical  and  valuable  of  modern 
scientific  marvels.  Often,  even  the  greatest  scientist  cannot  predict 
the  practical  value  of  theoretical  research. 





Hoyle,  Fred.  The  Nature  of  the  Universe.  Harper  and 
Brothers,  New  York,  1950.  This  is  a  very  small  book  that  gives 
an  exciting  picture  of  the  composition  of  the  stars  and  inter- 
stellar space. 

Jaffe,  Bernard.  Men  of  Science  in  America,  pp.  3S9-355. 
Simon  &  Schuster,  New  York,  1944.  Interesting  inioimation  on 
early  American  work  in  aeronautics. 

Walters,  Leslie.  "Chemistry  Exhibits  and  Projects."  Jour- 
nal of  Chemical  Education,  March,  1939,  pp.  113-115.  An 
illustrated  article  on  exhibits  and  projects  made  by  high- 
school  students.  Suggestions  for  what  you,  too,  could  do  along 
this  line. 


1.  A  reducing  agent  removes  oxygen  from  a  compound  con- 
taining oxygen.   (This  definition  will  be  somewhat  expanded 

2.  Reduction  and  oxidation  always  occur  in  the-  s.nnc  re- 

3.  Occlusion  is  the  absorption  of  gases  by  metals. 

4.  Careful  control  of  some  scientific  discoveries  is  necessary 
to  prevent  their  misuse. 

5.  Chemistry  needs  the  experimental  chemist,  the  scientist 
able  to  use  creative  imagination  in  formulating  theories,  and 
the  engineer  who  works  to  discover  how  theories  and  processes 
may  be  used  in  the  service  of  man. 


Group  A 

1.  (a)  Who  is  credited   with   the  discovery  of  pure  H2? 
(b)  How  was  H2  first  prepared? 

2.  (a)  What  substances  are  used  in  the  laboratory  prep- 
aration of  H2?  (b)  How  do  we  know  that  the  H..  comes  iroin 
the  acid  and  not  from  the  metal?  (c)  Write  the  word-equation 
for  this  reaction. 

3.  (a)  What  advantage  is  there  in  collecting  H2  by  the 
displacement  of  H2O?    (b)  What  kind  of  gas  could  not  be 
collected  in  this  manner? 


4.  (a)  Make  a  labeled  diagram  of  the  laboratory  prepara- 
tion of  H2.  (b)  Why  should  the  thistle  tube  extend  below  the 
surface  of  the  liquid  in  the  generating  bottle?   (c)  What  are 
two  reasons  for  using  a  thistle  tube? 

5.  (a)  When  Na  displaces  H2  from  H2O,  what  is  the  other 
substance  formed?  (b)  Why  can  you  not  see  it?  (c)  How  can 
you  obtain  it  for  inspection  of  its  properties? 

6.  What  metals  other  than  Na  are  so  active  that  they  will 
liberate  H2  from  water  when  the  metal  is  simply  placed  on 

7.  (a)  What  metal  will  liberate  H2  from  H2O  under  cer- 
tain special  conditions?    (b)  Describe  this  method  of  prepar- 
ing H2. 

8.  (a)  In  what  two  ways  is  HL,  prepared  tor  commercial 
use?    (b)  Why  were  these  methods  selected  in  preference  to 
ones  used  in  the  laboratory? 

9.  What  precaution  must  be  observed  in  preparing  and 
handling  H2? 

10.  (a)  List  five  physical  properties  of  H2.    (b)  Describe  a 
simple  experiment  to  illustrate  the  tact  that  H2  is  a  very  light 

11.  Determine  the  weight  and  cost  oi  the  H2  that  would  be 
needed  to  fill  a  dirigible  of  seven-million-cubic-foot  capacity. 
Consider  that  the  cost  of  the  H2  is  $2.00  per  hundred  cubic 
feet  of  gas. 

12.  What  element  in  the  air  is  used  in  the  burning  of  H2? 


13.  How  can  you  show  that  water  is  formed  when  hydrogen 

14.  What  is  reduction? 

15.  In  the  reduction  of  CuO  by  H2,  what  substance  is  oxi- 

16.  How  can  you  identify  each  of  the  products  obtained  in 
the  reduction  of  CuO  by  H2? 

17.  Oxidation   and  reduction   always  occur   together.   Ex- 

18.  (a)  Give  an  example  of  reduction  carried  out  on  a  large 
scale  in  industry,  (b)  What  reducing  agent  is  used? 


19.  (a)  What  is  the  function  of  each  part  of  a  bunsen 
burner?  (b)  What  is  meant  by  the  striking  back  of  the  flame? 

20.  (a)    What  conditions  are  necessary  for  a  flame?   (b) 
Make  a  labeled  drawing  of  the  flame  of  a  bunsen  burner. 

21.  What  formerly  useless  byproduct  is  converted  into  a 
very  useful  substance  by  the  use  of  H0? 

22.  How  is  oleomargarine  made? 

23.  In  what  special  fuels  is  H2  chiefly  responsible  for  the 
high  temperatures  obtained? 

Group  B 

24.  Can  you  suggest  a  sale  way  to  test  H,  for  its  burnability 
as  it  emerges  from  the  end  of  the  delivery  tube  in  the  labora- 
tory preparation? 

25.  Helium  gas  is  twice  as  dense  as  hydrogen  gas  and  yet 
can  lift  about  93%  as  much  weight  as  hydrogen  gas.  Explain. 

26.  (a)  Could  pure  H2  be  used  in  the  gas  range  at  home? 
(b)  Explain. 

27.  Why  does  pure  H2  burn  quietly  in  an  atmosphere  of  air, 
yet  burn  with  explosive  violence  when  the  two  gases  are  mixed 
and  ignited? 

28.  Tightly  bound  inflated  balloons  gradually  collapse.  Ex- 


1.  A  student's  mother  stopped  using  "Crisco"  after  her  son 
came  home  with  the  news  that  "Crisco"  was  not  a  "natural 
product"    but    was    manufactured    chemically,     (a)  Was    the 
mother   justified?     (b)  Give    reasons    for    your   answer    after 
you  have  consulted  your  family  physician,  the  producers  of 
"Crisco,"  and  the  U.S.  Department  of  Agriculture. 

2.  Organize  a  small  discussion  group  in  your  chemistry  class 
and  discuss  the  topic,  "Synthetic  chemistry  has  helped  in  the 
rise  of  totalitarian  states." 

3.  Prepare  a  report  or  organi/e  a  class  discussion  on  the 
topic   "Theoretical  science   has   irequently   resulted   in  great 
practical  discoveries." 

4.  WATER: 


.  .  .  Laboratories  are  necessary,  and, 
though  an  artist  without  a  studio  or 
an  evangelist  without  a  church  might 
conceivably  find  under  the  blue 
dome  of  heaven  a  substitute,  a  sci- 
entific man  without  a  laboratory  is  a 
misnomer.  Frederick  Soddy,  1877- 

Water  a  compound  of  two  gases  —  impossible!  For  thousands  of 
years,  water  was  considered  an  element.  Aristotle,  one  of  the  wisest 
of  Greeks,  included  this  liquid  among  the  "four  elements"  of  the 
ancients.  No  power  of  man  seemed  strong  enough  to  break  it  up 
into  any  recognizable  components.  Although  it  is  true  that  by  1780 
a  number  of  scientists  had  really  decomposed  water  so  that  hydrogen 
was  liberated,  they  were  unaware  of  the  real  nature  of  what  they 
had  done.  They  could  not  believe  that  hydrogen  had  actually  come 
from  the  water. 

"It  is  very  extraordinary  that  this  fact  should  have  hitherto  been 
overlooked  by  chemists.  Indeed,  it  strongly  proves  that  in  chemistry 
it  is  extremely  difficult  to  overcome  prejudices  imbibed  in  early 
education."  These  were  the  words  of  Lavoisier,  Often  it  is  hard  to 
overcome  superstition,  prejudice,  and  tradition.  However,  we  must 
learn  to  do  so  in  order  to  think  and  act  scientifically. 

In  1784,  Cavendish,  who  had  studied  hydrogen,  read  an  exciting 
paper  before  the  members  of  the  Royal  Society  of  England.  This 
is  what  he  told  them:  "Water  is  a  compound  of  oxygen  and  hydro- 
gen." What  a  startling  announcement!  Water  a  compound  of  two 
colorless,  tasteless  gases?  What  were  his  proofs?  Cavendish  told  them 




quietly  and  without  emotion.  He  said  that  he  had  placed  in  a  glass 
flask  a  mixture  composed  of  about  twice  as  much  air  as  hydrogen. 
Then  he  had  passed  an  electric  spark  through  the  mixture.  "All  the 
hydrogen  and  about  one-fifth  of  the  air  condensed  into  a  dew  which 
lined  the  glass.  In  short,"  he  continued,  "it  seemed  pure  water."  His 
experiments  proved  conclusively  that  water  is  a  compound  of  oxygen 
and  hydrogen,  and  yet  Cavendish  said  "it  seemed."  He  suspected  his 
listeners  would  not  be  convinced.  Water  a  compound  of  two  gases  — 

Lavoisier  convinces  the  world  that  water  is  a  compound.  Lavoisier, 
who  had  explained  the  nature  of  burning,  determined  to  tear  apart, 
or  analyze,  water  by  an  experiment  that  would  convince  the  world, 
just  as  Cavendish  had  shown  the  world  that  he  had  built  up,  or 
synthesized,  water  from  oxygen  and  hydrogen. 

He  arranged  the  apparatus  shown  below.  In  the  retort  A  he  heated 
pure  water,  so  that  steam  would  pass  through  the  tube  containing 
pure  charcoal,  which,  as  you  see,  was  heated  in  a  furnace.  The  gas 
that  escaped  passed  into  the  jar  H.  He  found  carbon  dioxide  gas 
dissolved  in  the  water  in  the  jar  and  identified  the  gas  issuing  from 
H  as  hydrogen.  The  water  (steam)  had  been  broken  up  into  hydro- 
gen, which  passed  on  as  a  gas,  and  oxygen,  which  combined  with  the 
carbon,  forming  carbon  dioxide.  Lavoisier  collected  and  weighed 
both  gases.  The  weight  of  all  the  resulting  products  accounted  for 
all  changes  in  the  weights  of  the  substances  used  in  the  reaction. 
Thus  Lavoisier  proved  without  a  doubt  that  water  could  be  broken 
up  into  the  two  gases  of  which  it  is  composed  —  hydrogen  and 

Fig.  13.  Apparatus  used  by  Lavoisier  to 
analyze  water.  Compare  this  equipment  with 
that  used  in  a  modern  laboratory. 

burning  charcoal 


hydrogen  gas 

reservoir  of  water  containing 
some  H2SO4 


Fig.  14.  Electrolysis  of  water.  Why  is  a  reservoir  necessary  in  this  apparatus? 

The  equation  for  the  chemical  changes  that  took  place  in  this 
experiment  is: 

2H2O  +  C  ->  2H2  +  CO2  (carbon  dioxide) 

How  water  can  be  broken  up  by  an  electric  current.  In  1800,  the 
invention  of  the  electric  battery  by  Volta  put  into  the  hands  of 
scientists  a  new  and  powerful  tool  for  determining  the  composition 
of  compounds.  The  decomposition  of  any  compound  by  electricity 
is  called  electrolysis. 

Within  a  few  months  following  Volta's  invention,  an  apparatus 
had  been  devised  for  the  electrolysis  of  water.  This  special  piece  of 
apparatus  is  used  in  the  laboratory  in  determining  the  composition 
of  water  by  volume.  Water  is  poured  into  the  center  glass  tube, 
which  acts  as  a  reservoir,  and  fills  the  other  two  tubes,  as  shown 
in  Fig.  14.  Since  pure  water  is  an  extremely  poor  conductor  of 
electricity,  that  is,  will  not  allow  electricity  to  pass  through  it  readily, 
a  small  amount  of  sulfuric  acid  is  added.  (At  the  close  of  the  chem- 
ical action,  the  amount  of  sulfuric  acid  is  unchanged.) 

The  electric  current  enters  the  apparatus  through  one  of  the 
platinum  electrodes,  passes  through  the  mixture  of  water  and  sulfuric 
acid,  and  leaves  the  apparatus  at  the  second  platinum  electrode. 
The  source  of  current  may  be  any  source  of  direct  current,  such  as 
a  storage  battery  or  dry  cells.  As  soon  as  the  circuit  is  closed,  bubbles 
of  gas  collect  on  the  surface  of  the  electrodes,  rise  through  the  water, 
and  collect  at  the  top  of  the  two  outside  tubes.  The  gas  that  collects 
at  the  negative  electrode,  or  cathode,  is  hydrogen;  the  gas  that  col- 
lects at  the  positive  electrode,  or  anode,  is  oxygen. 



The  composition  of  water  by  volume.  No  matter  when  we  stop 
the  action  of  the  current  in  the  electrolysis  of  water,  we  discover  a 
singular  fact.  The  volume  of  the  gas  at  the  cathode  is  twice  the 
volume  of  the  gas  at  the  anode.  In  other  words,  water  always  consists 
of  2  parts  of  hydrogen  to  1  part  of  oxygen  by  volume. 

If  the  gases  hydrogen  and  oxygen  are  mixed  in  a  closed  chamber 
and  then  exploded  by  an  electric  spark,  water  vapor  is  formed,  and 
the  ratio  in  which  these  gases  unite,  forming  water  vapor,  is  always 
2  volumes  of  hydrogen  to  1  of  oxygen.  Any  excess  of  either  gas  is  left 

The  composition  of  water  by  weight.  If  pure  water  is  thus  de- 
composed, forming  two  gases  in  a  definite  ratio  by  volume  (vol- 
umetric ratio) ,  and  if  water  is  formed  by  combining  these  gases  in 
the  same  volumetric  ratio,  it  ought  to  be  possible  to  find  in  this  com- 
bination a  constant  ratio  by  weight.  This  is  what  Cavendish  did  in 
1784  when  he  formed  water  by  combining  these  two  gases.  He  actu- 
ally weighed  both  the  gases  and  the  water.  Today  we  know  that  if 
we  combine  1  gram  of  hydrogen  (11  liters)  with  8  grams  of  oxygen 
(5.5  liters) ,  we  obtain  9  grams  of  water.  Thus,  the  composition  of 
water  by  weight  is  1  part  of  hydrogen  to  8  parts  of  oxygen. 

The  composition  of  water  by  weight  may  be  tested  in  the  labora- 
tory by  making  use  of  the  fact  that  pure  hydrogen  passed  over  heated 
copper  oxide  reduces  the  copper  oxide,  yielding  free  copper,  by 
uniting  with  the  oxygen  of  the  copper  oxide,  thus  forming  water 
(see  page  50)  .  The  weight  of  the  water  formed  is  always  found 
to  be  exactly  the  same  as  the  weight  of  the  hydrogen  used  plus  the 
weight  of  the  oxygen  lost  by  the  copper  oxide.  The  following  data 
are  the  results  of  such  an  experiment: 

Weight  of  copper  oxide  before  experiment 80  grams 

Weight  of  copper  left  after  experiment 64  grams 

Weight  of  oxygen  that  combined  with  hydrogen 16  grams 

Weight  of  water  produced 18  grams 

Weight  of  hydrogen  used 2  grams 

Result:  16  parts  by  weight  of  oxygen  combined  with  two  parts  of 
hydrogen  (or  8  parts  of  oxygen  combined  with  one  part  of  hydro- 
gen) ,  forming  water. 

.  induction 

latmum  electrodes  co;| 

Fig.  15.  Eudiometer.  In  this  appara- 
tus, a  mixture  of  hydrogen  and  oxy- 
gen is  ignited  by  an  electric  spark, 
forming  water.  Any  excess  of  either 
gas  acts  as  a  cushion  against  the 
rush  of  mercury  up  the  tube. 

Accuracy  and  satisfaction  in  scientific  investigations.  An  exact 
determination  of  the  relative  combining  weights  of  oxygen  and  hy- 
drogen was  essential  to  the  mathematics  of  chemistry.  Hundreds  of 
men  performed  thousands  of  experiments  to  determine  these  values 
as  accurately  as  human  ingenuity  could  devise.  Many  eminent  Ameri- 
cans were  among  them.  Edward  Morley  (1838-1923)  ,  a  professor  at 
Western  Reserve  University,  spent  more  than  ten  years  of  his  life  on 
such  experiments.  Finally  he  arrived  at  a  number  which  still  stands 
as  the  basis  for  chemical  calculations.  He  found  that  1.008  parts  of 
hydrogen  by  weight  combine  with  eight  parts  of  oxygen  by  weight, 
forming  water. 

Morley  and  the  others  did  not  receive  monetary  rewards  from 
these  experiments.  Theirs  was  a  labor  of  love,  a  work  of  pure  scien- 
tific research.  Their  sole  reward  was  the  satisfaction  of  knowing  that 
they  were  helping  to  increase  scientific  knowledge  whjch  might  be 
of  use  to  all  humanity.  The  world  needs  such  unselfish  men. 

The  law  of  definite  proportions.  You  have  learned  that  the  com- 
position of  water  is  always  the  same.  Perhaps  you  have  wondered 
if  this  is  a  property  of  all  chemical  compounds.  It  is  one  of  the 
fundamental  laws  of  chemistry  that,  in  forming  compounds,  elements 
combine  in  exact  proportions.  In  fact,  constancy  of  composition  is 
the  most  valid  test  used  in  deciding  whether  a  substance  is  a  com- 
pound or  a  mixture.  In  other  words,  the  composition  of  a  pure  com- 
pound never  varies.  This  is  the  law  of  definite  proportions.. 

"The  stones  and  soil  beneath  our  feet  and  the  ponderous  moun- 
tains are  not  mere  confused  masses  of  matter;  they  are  pervaded 
throughout  their  innermost  constitution  by  the  harmony  of  num- 
bers." This  is  indeed  fortunate,  for  if  the  composition  of  pure  com- 
pounds ever  changed,  instead  of  being  always  the  same,  the  exact 
measurements  that  we  use  in  chemistry  would  not  be  possible.  Par- 
ticularly is  this  true  of  quantitative  chemistry,  for  if  the  composition 
of  compounds  varied,  we  should  have  no  reliable  standard  chem- 
icals. Standards  are  necessary,  not  only  as  a  means  of  judging  the 
purity  of  a  substance  which  we  plan  to  use,  but  also  as  a  reliable 
basis  for  comparison. 


WATER  65 

What  is  a  "pure"  chemical?  Absolutely  pure  chemicals  are  almost 
impossible  to  make  or  buy.  Chemicals  that  have  no  appreciable  trace 
of  impurities  are  called  C.P.,  that  is,  chemically  pure.  The  designa- 
tion U.S.P.  refers  to  standards  of  purity  of  chemicals  to  be  used  in 
medicines  listed  in  the  United  States  Pharmacopoeia.  Such  chem- 
icals contain  no  harmful  impurities.  Reagent  grade  chemicals  con- 
form to  the  standards  of  the  American  Chemical  Society.  Tech  chem- 
icals do  not  meet  any  definite,  or  fixed,  standards  of  purity  but  are 
suitable  for  many  uses  in  which  slight  impurities  are  of  little  im- 

Patent  medicines,  other  remedies,  and  drugs,  such  as  aspirin,  are 
frequently  advertised  and  sold  under  trademarks.  The  purchaser 
should  always  carefully  examine  the  label  on  the  package  in  order 
to  be  sure  of  the  purity  of  the  contents.  The  chemical  composition 
of  the  contents  will  also  serve  as  a  guide  to  the  fair  price  of  the 
article.  Occasionally,  simple  and  inexpensive  C.P.  chemicals,  such 
as  bicarbonate  of  soda,  are  masked  behind  trademarks  and  sold  at 
exorbitant  prices.  The  consumer  can  help  to  guard  against  such 
practices  by  insisting  that  accurate  and  complete  statements  of  com- 
position be  printed  on  labels  of  all  packaged  goods. 

Some  physical  properties  of  water.  As  you  know,  water  is  an  odor- 
less, tasteless  liquid  that  is  colorless,  except  in  very  thick  layers,  when 
it  appears  blue.  Pure  water  freezes  at  0°C.  or  32°F.  and  boils  at 
100°C.  or  212°F.,  at  standard  conditions.  In  general,  impure  water 
has  a  higher  boiling  point  and  a  lower  freezing  point  than  pure 

The  fact  that  water  is  so  universally  distributed  has  led  to  its  use 
in  the  devising  of  scientific  standards  of  measure.  Thus  a  gram,  the 
metric  unit  of  weight,  is,  by  definition,  the  weight  of  a  milliliter  *  of 
chemically  pure  water  at  4°C.  This  temperature  is  chosen  because, 
when  water  cools,  it  contracts  until  the  temperature  reaches  4°C. 
Below  that  temperature  it  begins  to  expand  again.  Hence,  4°C.  is 
the  temperature  at  which  a  unit  volume  of  water  weighs  the  most. 
The  weight  of  a  unit  volume  of  a  substance*  is  known  as  its  density. 
Thus  at  4°C.  water  has  its  greatest  density.  Since  we  use  the  gram 
and  the  cubic  centimeter  or  the  milliliter  as  our  units  of  weight  and 
volume,  we  may  redefine  the  density  of  a  substance  as  the  weight  in 
grams  of  1  cubic  centimeter  or  1  milliliter  of  that  substance. 

*  A  milliliter  (ml.)  is  a  small  unit  of  capacity  (volume)  in  the  Metric  System. 
Numerically  it  is  equal  to  0.001  of  a  liter  and  is  used  in  measuring  the  volume  of 
fluids.  The  cubic  centimeter  (cc.)  is  also  used  in  measuring  volumes,  particularly 
of  solids. 







water  1.0 

Fig.  16.  The  relationship  of  specific  gravity  and 
buoyancy.  Aluminum  has  a  specific  gravity  greater 
than  that  of  water  and  does  not  float.  Ice  has  a  specific 
gravity  slightly  less  than  that  of  water  and  floats 
largely  submerged.  Cork  has  a  low  specific  gravity 
and  floats  with  most  of  its  mass  above  water. 

Since  water  has  a  density  of  7,  that  is,  1  milliliter  of  water  weighs 
1  gram  at  4°£.,  the  density  of  any  substance  is  also  the  ratio  of  the 
weight  of  a  given  volume  of  that  substance  to  the  weight  of  an  equal 
volume  of  water.  We  call  this  ratio  the  specific  gravity  (sp.  gr.)  of 
the  substance.  It  shows  the  comparison  between  the  weight  of  the 
substance  and  the  weight  of  an  equal  volume  of  water.  For  example, 
when  we  say  that  concentrated  su  If  uric  acid  has  a  specific  gravity 
of  1.84,  we  mean  that  it  is  1.84  times  as  heavy  as  water,  volume  for 
volume.  Since  below  4°C.  water  expands,  ice  is  lighter  than  water 
and  floats  on  it.  Ice,  of  course,  has  a  specific  gravity  less  than  that  of 




(In  grams  per  cubic  centimeter  or  per  milliliter) 

0.97  I  Iron  (pure)          7.86  |  Lead 



19.3     I  Platinum 


WATER  67 

The  high  specific  heat  of  water.  The  amount  of  heat  necessary  to 
raise  the  temperature  of  1  gram  of  water  1  degree  centigrade  is  called 
a  calory.  The  number  of  calories  necessary  to  raise  the  tempera- 
ture of  1  gram  of  a  substance  1  degree  centigrade  is  known  as  the 
specific  heat  (sp.  lit.)  of  that  substance.  Because  it  takes  1  calory  to 
raise  the  temperature  of  1  gram  of  water  1  degree  centigrade,  the 
specific  heat  of  water  is  1. 

Since  it  takes  only  one-thirtieth  as  much  heat  to  raise  the  tempera- 
ture of  1  gram  of  mercury  1  degree  centigrade  as  it  does  to  raise  1 
gram  of  water  1  degree  centigrade,  the  specific  heat  of  mercury  is  one- 
thirtieth  of  1,  or^j. 

Water  has  a  higher  specific  heat  than  most  other  substances.  Since 
it  requires  so  much  heat  to  raise  its  temperature,  it  warms  up  slowly. 
Conversely,  upon  cooling  it  gives  up  a  larger  amount  of  heat  for  the 
same  fall  in  temperature  than  most  other  substances  do.  It  is  partly 
because  of  the  high  specific  heat  of  water  that  it  is  used  in  home  heat- 
ing systems  and  in  the  cooling  systems  of  automobiles.* 

The  chemical  properties  of  water.  Water  is  a  stable  compound, 
that  is,  it  cannot  be  decomposed  easily.  It  does  not  even  begin  to 
decompose  into  hydrogen  and  oxygen  until  a  temperature  of 
1000°C.  is  reached.  Even  at  2500°C.  only  two  percent  of  it  is  decom- 
posed. However,  electricity,  in  the  presence  of  a  catalyst,  tears  it 
apart  easily  (see  page  62) . 

At  ordinary  temperatures,  water  is  decomposed  by  the  more  active 
metals,  such  as  sodium  and  potassium,  and  at  higher  temperatures 
by  the  less  active  metals,  such  as  iron.  In  these  cases  the  gas  liberated 
is  hydrogen.  Water  is  also  decomposed  by  the  more  active  nonmetals, 
such  as  chlorine  and  bromine,  but  these  liberate  oxygen  from  water 
instead  of  hydrogen. 

Water  acts  as  a  catalyst  in  many  chemical  reactions.  For  example, 
perfectly  dry  oxygen  and  hydrogen  do  not  unite  when  a  spark  is 
passed  through  them,  yet  the  faintest  trace  of  water  causes  such  a 
mixture  to  explode.  Phosphorus  does  not  burn  in  perfectly  dry  air, 
but  burns  readily  if  even  a  trace  of  water  vapor  is  present. 

*  Although  the  temperatures  of  boiling  water  and  steam  are  the  same,  it  takes 
about  540  calories  to  change  1  gram  of  water  from  its  boiling  point  of  100°C. 
to  steam  at  100°C.  This  amount  of  heat  is  called  the  heat  of  vaporization  of 
water.  Real  steam,  or  water  vapor,  the  gaseous  form  of  water,  is  invisible.  The 
visible  cloud  commonly  called  steam  is  water  vapor  after  it  has  condensed  into 
tiny  liquid  droplets. 

In  ice  water  the  temperatures  of  the  freezing  water  and  the  melting  ice  are 
both  0°C.  Yet  it  requires  about  80  calories  to  change  1  gram  of  ice  at  0°C.  to 
1  gram  of  water  at  0°C.  This  amount  of  heat  is  called  the  heat  of  fusion  of  ice. 

(left)  Photomicrograph  of  a  crystal  of  sodium  carbonate,  (right)  The  same  crys- 
tal after  a  few  hours  of  exposure  to  air.  What  has  occurred? 

Water  of  crystallization.  A  crystal  is  a  solid  mass  having  a  well- 
clefined  and  angular  form.  The  word  is  derived  from  a  Greek  word 
meaning  clear  ice.  Most  elements  and  compounds  are  capable  of  as- 
suming the  crystalline  form,  showing  sharp  edges  and  flat  surfaces. 
Such  a  substance  is  crystalline  washing  soda,  or  sodium  carbonate. 

When  a  crystal  of  washing  soda  is  heated  or  even  exposed  to  air, 
it  gives  off  water  and  crumbles  to  a  white  powder  which  is  not 
crystalline.  The  weight  of  water  liberated  bears  a  fixed  ratio  to  the 
weight  of  the  crystal  and  is  united  chemically  with  the  compound  of 
which  the  crystal  is  composed.  Water  which  is  thus  chemically  united 
with  a  substance  and  gives  that  substance  its  crystalline  form  is  called 
ivater  of  crystallization.  Such  water  is  rather  loosely  held  in  chemical 
combination  and  may  be  easily  expelled.  The  water  of  Crystallization 
is  separated  from  the  rest  of  the  formula  by  a  centered  dot,  which 
means  plus  (+)  and  is  not  a  multiplication  sign.  A  substance  that 
contains  water  of  crystallization  is  sometimes  called  a  hydrate. 

Another  common  hydrate  is  crystallized  copper  sulfate  once  com- 
monly known  as  blue  vitriol.  When  this  compound  is  heated,  its 
water  of  crystallization  is  liberated  and  it  crumbles  to  a  white  powder. 

CuS04  5H20       ->     CuS04        +        5H20 
crystallized  copper  sulfate  water  of 

copper  sulfate  (anhydrous)        crystallization 

This  change  in  color  is  further  evidence  that  the  water  of  crystalliza- 
tion is  chemically  united  with  the  copper  sulfate.  Use  is  made  of  the 
difference  in  color  between  white  anhydrous  copper  sulfate  and  the 
blue  hydrated  copper  sulfate  as  a  test  for  water.  Water  will  change 
anhydrous  copper  sulfate  to  the  blue  hydrate. 

The  ability  to  form  crystals  is  not  always  dependent  upon  the  pres- 
ence of  water.  Many  crystalline  substances,  such  as  table  salt  (so- 
dium chloride)  arid  sugar,  do  not  contain  water  of  crystallization. 
They  are  said  to  be  anhydrous,  meaning  without  water.  Crystals  that 
have  lost  their  water  of  crystallization  are  also  said  to  be  anhydrous. 




Efflorescent  substances  give  up  water.  Crystallized  washing  soda, 
on  exposure  to  air,  loses  its  water  of  crystallization  and  crumbles 
to  a  powder.  Such  a  substance  is  said  to  be  efflorescent,  which  means 
that  it  gives  up  its  water  of  crystallization  on  exposure  to  air.  The 
drier  the  air,  the  faster  the  loss  of  water  of  crystallization. 

Deliquescent  substances  take  up  water.  Dry  sodium  hydroxide, 
when  left  exposed  to  air,  soon  absorbs  enough  water  from  the  at- 
mosphere to  dissolve  itself  in  this  water.  Such  a  substance  is  said  to 
be  deliquescent.  The  higher  the  percentage  of  water  vapor  in  the 
air,  the  faster  the  process  of  deliquescence. 

Calcium  chloride,  a  white  solid,  is  deliquescent  and  is  often  used 
to  sprinkle  dry  roads  and  tennis  courts.  It  absorbs  moisture  from  the 
air  and,  in  this  way,  helps  to  keep  the  dust  down.  Magnesium  chlo- 
ride, an  impurity  found  in  common  table  salt,  is  also  deliquescent. 
Removal  of  the  magnesium  chloride  causes  pure  table  salt  to  remain 
dry  in  damp  weather,  to  pour  easily,  and  not  to  cake.  (What  sub- 
sance  does  your  mother  put  in  a  salt  shaker  to  keep  the  salt  from 
becoming  lumpy?) 

Deliquescent  substances  may  be  used  as  drying,  or  dehydrating, 
agents.  For  example,  concentrated  sulfuric  acid  absorbs  moisture 
from  the  air  and,  therefore,  is  used  in  drying  gases.  When  used  in 
the  laboratory,  these  dehydrating  agents  (the  most  common  of  which 
are  sodium  hydroxide,  sulfuric  acid,  and  calcium  chloride)  are  often 
placed  in  the  lower  compartment  of  a  vessel  known  as  a  desiccator; 
the  upper  compartment,  only  partially  separated  from  the  lower, 
contains  the  substance  to  be  dried. 

Our  lives  depend  on  water.  Water  is  essential  to  life.  Almost  70 
percent  of  the  total  weight  of  the  human  body  is  water,  and  plants 
contain  even  more.  Lettuce,  for  example,  contains  as  much  as  95 
percent  water  by  weight. 

,,..    .:-•  *v!-'"  -„  Fig.      IF*        Desiccator.     5yb» 

i.  *  to  be  are 

ln  the  VGpw  «m4  the 

is  put  In 



Immense  tracts  of  land  in  our  own  country,  such  as  the  hitherto 
arid  Columbia  Basin  of  the  Northwest,  have  been  or  will  be  turned 
into  rich  farmland  by  irrigation.  Federal  government  projects  have 
included  construction  of  huge  dams  such  as  the  Norris  Dam,  the 
Hoover  Dam,  the  Fort  Peck  Dam,  and  the  Grand  Coulee  Dam.  Be- 
hind mountainous  walls  of  earth  and  concrete  are  stored  huge  res- 
ervoirs of  water  which  are  changing  more  wasteland  into  fertile 
plains  and  are  helping  to  solve  the  problem  of  the  frequent  recur- 
rence of  disastrous  floods  in  certain  areas. 

Water  is  found  in  rocks,  paper,  fibers,  and  other  substances  gen- 
erally thought  of  as  "dry."  The  pages  of  this  book  may  contain  as 
much  as  10  percent  water  by  weight.  The  importance  of  water,  the 
most  common  solvent  in  nature,  is  discussed  in  detail  in  Chapter  15. 

Water  and  health.  Water  is  a  major  component  of  all  body  tissues 
and  fluids.  It  plays  a  major  role  in  the  preparation  of  foods  we  eat 
and  in  the  processes  of  digestion  and  assimilation.  Both  nutrients  and 
oxygen  are  carried  to  the  cells  of  the  body  in  fluids  composed  chiefly 
of  water,  and  many  of  the  waste  products  of  the  body  are  carried 
away  and  eliminated  in  a  similar  mariner.  \Vater,  in  the  form  of 
perspiration,  aids  in  regulating  the  temperature  of  the  body. 

To  maintain  normal  body  processes,  rather  large  quantities  of 
water  are  necessary.  The  amount  of  water  a  person  ^should  drink 
each  day  to  enable  these  processes  to  be  carried  on  varies  with  the 
kind  and  amount  of  activity,  the  temperature,  and  various  other 
factors.  However,  6  glasses  each  day  should  be  considered  a  minimum 
for  good  health. 

The  idea  that  water  should  not  be  drunk  with  meals  is  without 
foundation.  Digestion  proceeds  normally  even  when  large  quantities 
of  water  are  present  in  the  stomach.  However,  it  is  very  important 
not  to  substitute  the  drinking  of  water  with  meals  for  proper  and 
complete  chewing. 

This  illustration  shows  how  a  single  calcium  chloride  crystal  deliquesces,  gradually 
absorbing  water  from  the  air  until  it  becomes  completely  dissolved. 




Holmes,  Harry  N.  Out  of  the  Test  Tube  (4th  ed.) , 
Chap.  VII.  Emerson  Books,  New  York,  1945.  "The  Elixir  of 
Life,"  of  course,  refers  to  water. 

Read  the  Label  on  Foods,  Drugs,  Devices,  Cosmetics.  Cata- 
logue No.  FS  13.111:3/2,  1953.  Supt.  of  Documents,  Govt. 
Printing  Office,  Washington,  D.C.  Revision  No.  1  of  a  35-page 
illustrated  pamphlet  containing  valuable  information.  15^ 

Thorpe,  T.  E.  Essays  on  Historical  Chemistry,  pp.  98-122. 
The  Macmillan  Co.,  London,  1923.  A  description  of  the 
famous  controversy  (the  Water  Controversy)  over  the  priority 
of  the  discovery  of  the  composition  of  water,  involving  Cav- 
endish, Watt,  Priestley,  and  Lavoisier. 


1.  Analysis   is   the   breaking  down   of  a   compound   into 
simpler  substances. 

2.  Synthesis  is  the  building  up  of  a  more  complex  com- 
pound from  simpler  substances. 

3.  The  decomposition  of  any  compound  by  electricity  is 
called  electrolysis. 

4.  Edward  Morley  spent  more  than  ten  years  of  his  life  in 
determining  the  exact  ratio  in  which  oxygen  and  hydrogen 
unite  in  forming  water.  His  work  is  a  fine  example  of  ac- 
curacy and  patience  in  scientific  research. 

5.  Every  pure  chemical  compound  has  a  definite  composi- 
tion. This  is  the  law  of  definite  proportions.  C.P.  refers  to  a 
chemically  pure  substance. 

6.  The  density  of  any  substance  is  the  weight  of  a  unit 
volume  of  that  substance. 

7.  A  calory  is  the  amount  of  heat  necessary  to  raise  the 
temperature  of  one  gram  ol  water  one  degree  centigrade. 

8.  The  specific  heat  of  a  substance  is  the  number  of  calories 
necessary  to  raise  the  temperature  of  one  gram  of  that  sub- 
stance one  degree  centigrade. 

9.  The  specific  gravity  of  any  substance  is  the  weight  of 
one  cubic  centimeter  or  one  milliliter  of  that  substance  com- 
pared with  the  weight  of  an  equal  volume  of  water. 

10.  Water  of  crystallization  is  the  water  chemically  present 
in  certain  crystalline  substances. 

11.  An  efflorescent  substance  loses  water  of  crystallization 
on  exposure  to  air;  a  deliquescent  substance  absorbs  water 
from  the  air. 



Group  A 

1.  Why  is  water  considered  a  compound? 

2.  How  did  Cavendish  synthesize  H2O? 

3.  What  is  the  difference  between  analysis  and  synthesis? 

4.  Briefly  describe  how  Lavoisier  analyzed  H2O. 

5.  What  is  electrolysis? 

6.  (a)  Make  a  labeled  diagram  of  the  apparatus  for  the 
laboratory  electrolysis  of  H2O.  Indicate  the  direction  of  the 
current,  the  cathode,  and  the  anode,  (b)  At  which  electrode  is 
the  H2  given  off?  (c)  the  O2? 

7.  What  part  does  sulfuric  acid   (H2SO4)   play  in  the  elec- 
trolysis of  H2O? 

8.  Write  the  word-equation  for  the  electrolysis  of  H2O. 

9.  How  would  you   test  to  find  out  which  of  the  gases 
present  in  the  two  outer  tubes  of  the  electrolysis  apparatus 
is  H2? 

10.  (a)  What  is  the  composition  of  H2O  by  volume?  (b)  by 

11.  What  are  the  differences  between  C.P.,  V.S.P.,  Tech, 
and  Reagent  chemicals? 

12.  In  connection  with  the  study  of  H2O,  cite  an  example  of 
the  accuracy  and  patience  ol  men  of  science. 

13.  State  and  illustrate  the  law  of  definite  proportions. 

14.  (a)  Does  vigorously  boiling  H2O  have  a  higher  tempera- 
ture than  slowly  boiling  H2O?    (b)  Explain. 

15.  What  are  five  physical  properties  of  H2O? 

16.  How  does  steam  differ  from  water  vapor? 

17.  Why  is  it  sometimes  unsafe  to  purchase  drugs  or  medi- 
cines by  trade-mark  alone? 

18.  Explain  why  water  pipes  often  burst  in  cold  weather. 

19.  (a)  Which  would  you  prefer  to  heat  your  feet  on  a  cold 
night  — a  hot  flatiron  or  a  bottle  of  hot  H2O?    (b)  Explain. 

20.  If  H2O  were  less  stable  than  it  is,  what  disaster  might 

WATER  73 

21.  Describe  two  chemical  properties  of  H0O. 

22.  Give  an  example  of  the  part  that  a  trace  of  H2O  may 
play  in  bringing  about  a  chemical  change. 

23.  Illustrate  what  is  meant  by  water  of  crystallization. 

24.  Transparent  crystals  of  washing  soda  change  to  a  white 
powder.  Is  this  a  physical  or  a  chemical  change? 

25.  Why    is    an    unstoppered    bottle    of    calcium    chloride 
(CaCl2)  sometimes  left  in  large  clocks? 

26.  When  crystals  of  table  salt    (NaCl)    are  heated,  some 
H2O  is  liberated.  Is  this  water  of  crystallization?  Explain. 

27.  Examine  some  pellets  of  NaOH  that  have  been  exposed 
to  air.  What  property  does  NaOH  have? 

28.  (a)  Which  gives  a  more  severe  burn,  boiling  H2O  or 
steam?  (b)  Why? 

29.  Explain  the  importance  ol  H2O  to  health. 

Group   B 

30.  If  H2O  did  not  expand  on  Iree/ing,  how  would  it  affect 

31.  A  spark  is  passed  through  a  mixture  ol  60  ml.  of  O2  and 
50  ml.  of  H2  in  the  presence  of  water  vapor.  What  substances, 
and  how  much  of  each,  will  be  found  in  the  tube  after  the 

32.  Devise  an  experiment  to  show  the  composition  of  H.,O 
by  weight.  Mention  actual  weights. 

33.  How   would    you    determine    the    percentage    of    H.,O 
present  in  a  sample  of  "dry"  wood? 

34.  A  liter  of  H0  weighs  0.09  g.  and  a  liter  of  O2  weighs 
1.43  g.  Show  how  you  could  find  the  composition  of  H2O  by 
weight  from  these  facts  (data) . 

35.  What  weight  of  oxygen  can  be  obtained  from  the  elec- 
trolysis of  50  pounds  of  water? 

36.  Ice  is  purer  than  water.  Would  it  be  safe  to  use  ice  from 
a  polluted  pond  in  your  iced  tea?  Explain. 


1.  Find  out  from  your  mother  or  grocer  the  cost  of  pound 
packages  of  powdered  *washing  soda  and  crystallized  washing 


soda.  The  water  of  crystallization  present  in  crystallized  wash- 
ing soda  is  equal  to  approximately  63  percent  of  the  weight 
of  the  crystal.  On  this  basis,  calculate  which  is  less  expensive, 
the  crystallized  washing  soda  containing  water  of  crystallization 
or  the  powdered  washing  soda  containing  no  water  of  crystal- 
lization. Make  a  report  of  your  problem  in  class. 

2.  A  student  came  to  his  chemistry  teacher  very  excited 
about  his  invention  of  a  system  that  would  run  an  airplane 
indefinitely  on  a  small  amount  of  H2O.  He  planned  to  de- 
compose H2O  by  electricity,  use  the  O2  and  H2  thus  produced 
to  supply  an  oxyhydrogen  torch  which,  when  ignited,  would 
boil  H2O  for  a  steam  engine.  The  steam  engine  in  operation 
would  give  power  to  the  plane's  propeller  and,  at  the  same 
time,  to  an  electric  generator.  The  electricity  thus  produced 
would  be  harnessed  to  decompose  more  H2O,  which  would  be 
constantly  renewed  as  a  product  of  the  burning  of  the  O2  and 
H2  in  the  oxyhydrogen  torch.  What  do  you  think  of  this  in- 
vention?  Your   answer   should   be   complete   with   scientific 

3.  Benjamin  Thompson,  one  of  the  greatest  of  early  Amer- 
ican scientists  and  who  later  became  Count  Rumford,  made 
classic  contributions  in  the  field  of  specific  heat.  "It  is  a  strange 
coincidence,"  wrote  Albert  Einstein,  "that  nearly  all  the  funda- 
mental work  concerned  with  the  nature  of  heat  was  done  not 
by  professional  scientists  but  by  men  who  regarded  science  as 
their  great  hobby."  Mention  two  other  scientific  contributions 
made  by  men  who  were  not  professional  scientists. 

4.  Get  some  anhydrous  copper  sulfate  and  with  it  determine 
whether  certain  "dry  substances"  really  contain  water.  Report 
your  results. 


The  bodies  which  time  and  nature 
add  to  things  little  by  little,  .  .  . 
no  exertion  of  the  eyesight  can  be- 
hold; and  so,  too,  wherever  things 
grow  old  by  age  and  decay,  and  when 
rocks  hanging  over  the  sea  are  eaten 
away  by  the  gnawing  salt  spray,  you 
cannot  see  what  they  lose  at  any 
given  moment.  Nature,  therefore, 
works  by  unseen  bodies. 

Lucretius,  99-55  B.C. 

The  origin  of  the  idea  of  atoms.  The  most  fruitful  scientific  specu- 
lation that  came  out  of  ancient  Greece  2500  years  ago  was  that  mat- 
ter is  made  up  of  small,  eternal  particles  in  continual  motion.  Leu- 
cippus  and  his  pupil  Democritus  (de  mok'ri  tws) ,  taught  that  all 
matter  was  composed  of  invisible,  indivisible,  indestructible  parti- 
cles, or  atoms.  During  the  seventeenth  century,  Newton  had  similar 
ideas  about  the  nature  of  matter.  "It  seems  probable  to  me,"  wrote 
Newton,  "that  God  in  the  beginning  formed  matter  in  solid,  massy, 
hard,  impenetrable,  moveable  particles,  so  very  hard  as  never  to  wear 
or  break  into  pieces."  The  gradual  development  of  the  idea  of  the 
atom  is  an  interesting  story. 

How  Dalton's  approach  to  the  nature  of  matter  differed  from  that 
of  the  Greeks.  Early  in  the  nineteenth  century,  an  English  scientist, 
John  Dalton,  became  interested  in  the  idea  of  atoms.  Newton's  idea 
was  a  beautiful  one,  thought  Dalton,  but  did  it  check  with  the  known 
facts?  Would  it  help  explain  some  of  the  physical  properties  of  gases, 
which  had  so  puzzled  him? 

It  is  interesting  to  note  the  difference  in  the  way  the  Greek  teach- 
ers had  come  to  their  ideas  and  the  way  in  which  Dalton  formulated 
his  theory.  The  Greek  teachers  made  a  few  observations,  followed. 




some  logical  reasoning,  and  then  ventured  an  opinion.  For  example, 
they  noticed  that  a  lump  of  salt  could  be  broken  down  into  bits  of 
salt,  which  could  then  be  further  reduced  to  tiny  crystals  of  salt.  It 
was  inconceivable,  they  reasoned,  to  continue  this  division  forever. 
There  must  come  a  time,  therefore,  when  one  would  finally  reach  the 
smallest  piece  of  salt,  that  is,  an  indivisible  atom  of  salt.  Dal  ton  did 
more.  He  also  experimented.  He  tried  to  find  out  why  the  gases  of 
the  atmosphere  remained  mixed,  how  gases  dissolve  in  water,  and 
whether  the  composition  of  pure  compounds  varies  or  is  constant. 
On  the  basis  of  some  not  too  accurate  observations  in  these  experi- 
ments and  some  indirect  evidence  from  facts  known  in  his  day, 
he  formulated  the  atomic  theory.  In  1803  he  announced  it  suddenly 
without  waiting  to  test  all  of  it  by  experiment.  Since  that  time  a 
great  many  additional  experiments  have  proved  that  Dalton's  theory 
was  substantially  correct.  ,  ; 

With  this  electron  microscope, 
scientists  may  study  molecules 
and  other  infinitesimal  par- 
ticles. Typical  photomicro- 
graphs may  be  seen  in  the 

ATOMS  77 

Dalton's  atomic  theory.  Dalton's  atomic  theory  was  based  on  the 
following  ideas: 

1)  That  all  matter  consists  of  extremely  small  particles,  called 

2)  That  all  the  atoms  of  any  one  element  are  alike  in  size,  shape, 
and  weight. 

3)  That  the  atoms  of  one  element  differ  from  those  of  all  other 
elements  in  size,  shape,  and  weight. 

4)  That  chemical  changes  are  changes  in  the  combination  of  atoms 
with  each  other. 

5)  That  even  in  the  most  violent  chemical  change,  the  indivisible 
atoms  do  not  break  into  pieces. 

The  diameter  of  the  hydrogen  atom  is  about  1/250,000,000  inches. 
This  is  several  hundred  times  smaller  than  the  average-sized  bac- 

To  explain  his  theory,  Dalton  drew  pictures  of  atoms.  Each  atom 
was  represented  by  a  circle.  Since  the  atoms  of  elements  are  unlike, 
he  varied  the  pictures  of  the  circles  as  follows: 

Hydrogen  Oxygen  Nitrogen 

Carbon  0  Sulfur  Q  Gold 

Dalton  pictured  chemical  change  as  the  union  of  one  or  more 
atoms  of  one  element  with  atoms  of  other  elements.  He  believed  that 
when  mercury  is  heated  in  air,  one  atom  of  mercury  unites  with  one 
atom  of  oxygen,  forming  a  particle  of  the  compound,  mercuric  oxide. 
To  demonstrate  this  union  of  atoms,  Dalton  constructed  model 
spheres,  bringing  them  into  contact  with  each  other: 

1  atom  of  liquid         1  atom  of  gaseous  1  molecule  of  red 

mercury  "*"  oxygen  *      mercuric  oxide 

According  to  Dalton,  atoms  preserve  their  individuality  in  all 
chemical  changes.  Hence,  Dalton  described  an  atom  as  the  smallest 
part  of  an  element  that  takes  part  in  a  chemical  change  without  itself 
being  altered.  Atoms  combine  to  form  molecules.  Two  or  more  atoms 
may  combine  into  a  molecule  of  an  element  or  into  a  molecule  of  a 
compound.  A  molecule  is  the  smallest  part  of  a  compound  or  element 
that  has  the  chemical  properties  of  that  compound  or  element. 


Inertia  in  scientific  thinking.  Dalton's  theory  was  strongly  at- 
tacked by  the  leading  scientists  of  his  day.  One  of  the  most  eminent 
of  them  said  he  could  not  understand  "how  any  man  of  sense  or  sci- 
ence would  be  taken  in  by  such  a  tissue  of  absurdities."  Dalton's 
theory  was  the  result  of  creative  imagination  and  the  boldness  of  a 
great  thinker.  Dalton  had  never  seen  nor  weighed  an  atom.  Yet  his 
theory  was  of  practical  value  and  was  accepted  gradually  by  the  scien- 
tific world  as  a  useful  working  hypothesis  by  which  chemical  changes 
could  be  explained. 

The  newer  electron  theory  of  matter  (see  Chapter  11)  has  ex- 
panded Dalton's  theory.  It  has  been  modified  in  details,  but  its  gen- 
eral applications  still  hold. 

The  use  of  Dalton's  theory.  One  of  the  questions  that  scientists  of 
Dalton's  time  were  studying  was  whether  the  composition  of  a  com- 
pound is  always  the  same  or  whether  it  varies.  Some  believed  that 
compounds  are  always  formed  from  fixed  amounts  of  elements.  They 
believed  therefore,  that  the  composition  of  a  compound  is  always  the 

The  French  chemist,  Claude  Berthollet  (ber'to'le')  ran  some  ex- 
periments to  test  this  question.  On  the  basis  of  his  experiments  he  be- 
lieved that  the  composition  of  compounds  might  vary  to  some  extent. 

Joseph  Proust  (prdost) ,  another  Frenchman,  set  out  to  settle  this 
difference  of  opinion.  He  repeated  Berthollet's  experiments,  using 
the  purest  of  chemicals  and  the  most  delicate  apparatus  available. 
Taking  every  precaution  to  prevent  error,  he  found  mistakes  in 
Berthollet's  work.  He  found  that  his  fellow-scientist  had  used  im- 
pure compounds  and  substances  such  as  glass  and  mixtures  of  metals 
(alloys)  and  mixtures  of  liquids,  which  were  not  pure  compounds. 

For  eight  long  years,  the  difference  of  opinion  persisted.  Never, 
however,  did  it  become  anything  but  an  honest,  truth-seeking  discus- 
sion. Personal  whims  and  prejudices  did  not  decide  the  matter. 
When  Berthollet  considered  Proust's  evidence  and  discovered  his 
own  errors,  he  accepted  Proust's  verdict  and  agreed  that  the  compo- 
sition of  compounds  is  always  the  same. 

The  law  of  definite  proportions.  This  law  states  that  the  elements 
in  a  compound  always  occur  in  a  definite  proportion  by  weight.  This 
is  another  way  of  saying  that  the  composition  of  compounds  is  always 
the  same. 

Dalton's  little  circles  very  neatly  explained  the  law.  The  weight 
of  atoms  of  any  element  is  always  the  same.  Compounds  are  com- 
posed of  these  minute  and  unchangeable  atoms.  Therefore  the  com- 
position of  compounds  by  weight  must  be  definite  and  uniform. 



Dalton  discovers  the  law  of  multiple  proportions.  Dalton  knew 
that  one  atom  of  carbon  •  unites  with  one  atom  of  oxygen  O  to 
produce  the  deadly  gas,  carbon  monoxide  •O  •  In  this  compound 
the  carbon  weighs  f  as  much  as  the  oxygen.  This  fraction  can  be  ex- 
pressed as  the  ratio  of  3  to  4. 

Carbon  also  combines  with  oxygen  to  form  carbon  dioxide.  Dal- 
ton wrote  this  combination  as  0QO-  In  this  compound  the  carbon 
weighs  only  f  as  much  as  the  oxygen.  The  ratio  is  three  parts  of  car- 
bon to  eight  parts  of  oxygen:  3  to  8.  From  observing  this  and  other 
similar  combinations,  Dalton  formulated  another  fundamental  law 
of  chemistry,  the  law  of  multiple  proportions. 

You  note  that  in  both  carbon  monoxide,  CO,  and  in  carbon  di- 
oxide, CO2,  the  weight  of  the  carbon  is  the  same.  But  the  weight  of 
the  oxygen  in  carbon  monoxide  is  4,  and  in  carbon  dioxide  is  8. 
Thus  three  parts  of  carbon  combine  with  either  4  parts  or  8  parts  of 

When  two  elements  combine  to  form  more  than  one  compound, 
with  the  weight  of  one  element  remaining  fixed,  the  ratios  of  the 
weights  of  the  other  elements  are  small  whole  numbers. 

Thus  the  amounts  of  oxygen  that  unite  with  three  parts  of  carbon 
are  in  the  ratio  of  4  to  8,  or  1  to  2. 

How  the  discovery  of  hydrogen  peroxide  helped  to  uphold  the  law 
of  multiple  proportions.  In  1818,  Louis  Thenard  (ta  nar') ,  a  French 
teacher  of  chemistry,  discovered  a  compound,  which  upon  analysis 
was  shown  to  be  made  up  of  equal  volumes  of  oxygen  and  hydrogen. 
This  compound  is  hydrogen  peroxide. 

Hydrogen  and  oxygen  combine  to  form  two  different  compounds. 
Water  is  composed  of  one  part  hydrogen  and  eight  parts  oxygen  by 
weight.  Hydrogen  peroxide  is  composed  of  one  part  hydrogen  and 
sixteen  parts  oxygen.  Thus  the  ratio  of  the  weights  of  oxygen  that 
combine  with  a  fixed  weight  of  hydrogen  is  8  to  16,  or  1  to  2. 





(  Water  H2O 
(^  Hydrogen  peroxide  H2O2 

1  H 
1  H 

80  "I 

160  / 

1  to  2 

f  Carbon  monoxide  CO 
^  Carbon  dioxide  CO2 



8O  ) 

1  to  2 

This  maze  of  tanks  and  pipes  is  re- 
quired for  the  commercial  preparation 
of  hydrogen  peroxide. 

Properties  and  uses  of  hydrogen  peroxide,  FLO,.  Since  water  and 
hydrogen  peroxide  have  different  chemical  compositions,  they  have 
different  physical  and  chemical  properties.  Hydrogen  peroxide  is  a 
colorless  liquid,  about  one  and  one-half  times  as  heavy  as  water.  It 
is  odorless  and  mixes  with  water,  alcohol,  or  ether.  It  is  useful  com- 
mercially because  it  is  unstable.  That  is,  heat  or  light  decomposes  it 
easily  into  water  and  oxygen: 

H2O2  -*  H20  +  O  T 

The  oxygen  atom  that  is  liberated  is  in  a  very  active  state,  i^eady  to 
combine  with  another  atom  of  oxygen  or  with  any  other  substance 
at  the  instant  of  liberation.  This  very  active  atomic  oxygen  is  some- 
times called  nascent  (newborn) .  Nascent  oxygen  is  written  as  (X 
ordinary  oxygen  as  O2.  The  arrow  pointing  upwards  represents  a  gas. 

Some  colored  compounds  lose  their  color  when  they  are  oxidized. 
Fibers  containing  compounds  which  give  them  their  color  can  be 
bleached  by  exposing  them  to  nascent  oxygen.  Hydrogen  peroxide 
is  used  as  an  oxidizing  agent  to  bleach,  or  decolorize,  cotton  goods, 
wool,  wood  pulp,  wood  used  for  furniture,  as  well  as  silk,  hair,  feath- 
ers, glue,  and  other  animal  substances. 

Some  bacteria  are  destroyed  when  exposed  to  oxygen.  Hydrogen 
peroxide  is  therefore  used  as  a  household  antiseptic.  The  household 
product  actually  is  mainly  water  with  a  small  amount  (usually  3  per 
cent)  of  hydrogen  peroxide  dissolved  in  it.  It  also  contains  some  sub- 
stance such  as  acetanilid,  which  retards  the  decomposition  of  the  hy- 
drogen peroxide.  As  an  antiseptic,  hydrogen  peroxide  is  not  very 
effective  since  the  oxygen  it  releases  does  not  reach  enough  of  the 




Commercial  preparation  of  hydrogen  peroxide.  It  barium  perox- 
ide, a  white  solid,  is  treated  with  dilute  sulfuric  acid  at  a  temperature 
below  Lr)°C.,  hydrogen  peroxide  and  barium  sultate  are  formed.  Bar- 
ium sulfate  is  a  white  insoluble  solid  which  settles  out. 

Sulfuric  acid  -f  barium  peroxide  —  >  hydrogen  peroxide  4-  barium  sulfate 
;H2!SO4~~~+"~~~BaO2j          ->  H2O2  +        BaSO4  1 

This  type  of  reaction  is  called  double  replacement  —  barium  replaces 
hydrogen.  The  arrow  pointing  downward  after  BaSO4  indicates  that 
this  compound  is  insoluble  and  separates  out,  or  precipitates.  The 
insoluble  substance  that  separates  out  is  called  a  precipitate. 

Most  of  the  hydrogen  peroxide  produced  commercially  is  made  by 
gently  heating  pcrsulfuric  acid,  H_,S,ON,  which  reacts  with  water: 

Persulfuric  acid 




sulfuric  acid 

+     hydrogen  peroxide 
+  H2O2 

In  this  process  the  hydrogen  peroxide  is  distilled  out.  Superoxol  is  a 
30%  H2O2  solution.  A  90%  solution  of  this  substance  is  used  as  a 
rocket  fuel. 

The  Glenn  Martin  Company 

Off  to  an  altitude  of  160  miles  roars 
the  Navy's  Viking  No.  1 1  rocket. 
In  flight  this  rocket  will  attain  a 
speed  of  4300  miles  per  hour. 
Rocket  development  depends 
greatly  upon  fuel  research  carried 
on  in  chemical  laboratories. 



Harrow,  Benjamin.  The  Romance  of  the  Atom,  pp.  27-33. 
Boni  &  Liveright,  New  York,  1927.  Origin  and  development 
of  ideas  about  atoms. 

Langdon-Davies,  John.  Inside  the  Atom.  Harper  &  Bros., 
New  York,  1933.  Amusing,  popular  introduction  to  science 
and  the  nature  of  matter. 

Leicester,  H.  M.  and  Klickstein,  H.  S.  Source  Book  in  Chem- 
istry, pp.  215-220.  McGraw-Hill  Book  Company,  New  York, 
1952.  Gives  John  Dal  ton's  observations  on  the  constitution  of 

Thomson,  J.  Arthur.  The  Outline  of  Science,  pp.  245-253. 
G.  P.  Putnam's  Sons,  New  York,  1937.  "Foundations  of  the 
Universe."  In  "The  World  of  Atoms"  the  size  and  energy  of 
these  tiny  particles  are  very  simply  discussed. 


1.  The  idea  of  individual  indivisible  particles  of  matter 
originated  with  the  ancient  Greek  teachers.  Dalton  used  this 

[idea  and  developed  the  atomic  theory  in  1803. 

2.  The  chief  assumptions  of  Dalton's  theory  were:  a)  That 
all  matter  consists  of  extremely  small  particles,  called  atoms, 
b)  That  all  the  atoms  of  any  one  element  are  alike  in  size, 
shape,  and  weight,  c)  That  the  atoms  of  one  element  differ 
from  those  of  all  other  elements  in  size,  shape,  and  weight, 
d)  That  chemical  changes  are  changes  in  the  combination  of 
atoms  with  each  other,   e)  That  even  in   the  most  violent 
chemical  changes,  the  indivisible  atoms  do  not  break  into 

3.  An  atom  is  the  smallest  part  of  an  element  that  takes 
part  in  a  chemical  change  without  itself  being  altered. 

4.  A  molecule  is  composed  of  two  or  more  atoms  and  is  the 
smallest  part  of  a  compound  or  element  that  has  the  proper- 
ties of  that  compound  or  element. 

5.  Personal  whims,  prejudices,  or  prestige  should  play  no 
part  in  settling  scientific  problems. 

6.  The  law  of  definite  proportions  states  that  the  elements 
in  a  compound  always  occur  in  a  definite  proportion  by 

7.  The  law  of  multiple  proportions  states:  When  any  two 
elements  combine  to  form  more  than  one  compound,  with  the 
weight  of  one  element  remaining  fixed,  the  ratios  of  the  weights 
of  the  other  element  are  small  whole  numbers. 




Group  A 

1.  When  did  John  Dalton  advance  his  atomic  theory? 

2.  Was  the  idea  of  atoms  new  in  Dalton's  time?  Explain. 

3.  How  did  Dalton's  approach  to  the  study  of  the  nature  of 
matter  differ  from  that  of  the  Greeks? 

4.  State  the  five  essentials  of  Dalton's  atomic  theory. 

5.  How  did  Dalton  distinguish  between  atoms  of  different 

6.  According   to   Dalton,   what   happens    when   elements 

7.  How  did  the  scientific  world  receive  Dalton's  theory? 

8.  Describe  the  difference  of  opinion  between  Berthollet 
and  Proust. 

9.  (a)    State  the  fundamental  law  of  chemistry  that  Dalton 

(b)  Show  how  it  applies  to  the  composition  of  CO  and 

10.  (a)  When  and  (b)  by  whom  was  H2O2  prepared  for  the 
first  time? 

1 1 .  Compare  the  physical  properties  of  H2O  and  H2O2. 

12.  Contrast  the  chemical  properties  of  H2O  and  H2O2. 

13.  Compare  the  compositions  of  H2O  and  H2O2  by  volume 
and  weight. 

14.  State  two  ways  in  which  the  decomposition  of  H2O2  may 
be  retarded. 

15.  Write  the  word-equation  for  the  decomposition  of  H2O2. 

16.  What  is  the  difference  between  O2  and  nascent  oxygen? 

t  .  .  . 

17.  Describe  the  two  most  useful  properties  of  H2O2. 

18.  What  is  the  great  advantage  of  H2O2  as  a  bleaching 

19.  What  is  an  antiseptic? 

20.  State  a  commercial  method  of  preparing  H2O2. 

21.  Write  the  equation  for  the  preparation  of  H2O2  from 


22.  What  is  a  precipitate? 

23.  Compare  the  chemical  properties  of  O2,  O3,  and  O. 

Group  B 

24.  Explain  how  the  law  of  multiple  proportions  is  based 
on  Dalton's  atomic  theory. 

25.  Two  oxides  of  nitrogen,  nitrous  oxide    (N.,O)    and  ni- 
trogen dioxide    (NO2) ,  show  the  following  ratios  by  weight: 
In  nitrous  oxide,  the  ratio  by  weight  of  nitrogen  to  oxygen 
is  7  to  4;  in  nitrogen  dioxide,  the  ratio  by  weight  of  nitrogen 
to  oxygen  is  14  to  32.  With  what  fundamental  law  of  chemistry 
are  these  figures  in  accord? 

26.  Do  we  still  believe  that  all  atoms  of  the  same  element 
weigh  the  same? 

27.  (a)  Is  Dalton's  atomic  theory  still  a  theory  or  has  it  been 
proved  experimentally?  (b)  Explain  your  answer. 

28.  Dalton's   hobby   of   recording  weather   conditions   was 
greatly  responsible  for  the  atomic  theory.  Can  you  cite  another 
example  in  science  where  a  hobby  has  resulted  in  a  great 

29.  Write  a  2-  or  3-page  report  on  the  life  of  John  Dalton. 


1.  Dalton's  atomic  theory  is  a  beautiful  example  of  creative 
imagination  in  pure  science.  Until  very  recently  America,  al- 
though it  excelled  Europe  in  inventions  and  applied  science, 
has  lagged  behind  the  Old  World  in  the  kind  of  creative  imagi- 
nation represented  by  Dalton's  theory  and  the  germ  theory  of 
disease  which  Pasteur  gave  humanity.  Can  you  give  reasons 
for  this  state  of  affairs?  What  of  the  future? 

2.  Add  a  pinch  of  manganese  dioxide  to  one-third  of  a  test 
tube  of  hydrogen  peroxide  from  your  medicine  cabinet.  Insert 
a  glowing  splint.  Prepare  a  brief  report  on  the  reaction  which 
takes  place. 

3.  Prepare  a  debate  or  a  class  discussion  on  the  topic,  "If 
prejudices  and  superstitions  did  not  exist,  social  change  would 
proceed  more  rapidly." 

4.  Form  a  committee  to  investigate  and  report  on  prejudices 
and  superstitions  common  to  your  community.  Do  any  of  these 
have  a  scientific  foundation? 

5.  Make  models  of  the  atoms  of  hydrogen,  carbon,  sulfur, 
gold,  mercury,  and  nitrogen  as  Dalton  constructed  them,  using 
marbles  or  some  other  suitable  spheres. 




The  mysterious  symbols  of  alchemy.  Ancient  alchemists  feared 
that  they  would  lose  their  unique  social  position  if  anyone  could 
understand  their  work.  Consequently  they  used  strange  symbols, 
both  to  conceal  the  true  nature  of  their  writings  and  also  to  give 
themselves  an  added  air  of  mystery  and  magic.  They  represented 

sand  by     .jjj.^  ,  glass  by  Q—  Q,  soap  by  ^^-,  and  salt  by  ©.  To 
•  •  •  •  •  ^^^^^  ^"^ 

them  the  symbcJ  for  perfection  and  also  for  the  sun  was  the  circle 
O  •  Hence  it  was  used  to  represent  sun-colored  gold,  which  they  con- 
sidered the  perfect  metal. 

Many  of  their  symbols  were  derived  from  ancient  mythology.  The 
lance  and  shield  of  Mars,  god  of  war,  £f  ,  represented,  appropriately 

enough,  iron.  The  looking  glass  of  Venus,  O,  was  their  symbol  for 
copper  because,  according  to  legend,  Venus  had  first  appeared  on  the 
shores  of  the  island  of  Cyprus,  long  famous  for  its  copper  mines. 

Often  an  alchemist  would  develop  a  set  of  symbols  for  his  own 
special  use  and  would  reveal  the  meaning  to  no  one,  not  even  to  his 
brother  alchemists.  In  one  Italian  manuscript  written  in  the  seven- 
teenth century,  the  element  mercury  was  represented  by  20  different 
symbols  and  35  different  names.  As  long  as  alchemy  was  a  purely 
personal  practice  carried  on  for  selfish  ends,  the  confusing  symbolism 
presented  no  problem.  However,  when  the  universal  science  of 
modern  chemistry  began  to  emerge,  it  was  essential  to  develop  a 
system  of  symbols  which  could  be  understood  easily  in  every  country 
of  the  world.  Dalton,  and  a  few  others  before  him,  had  attempted 
to  substitute  some  reasonable  system  for  the  jungle  of  weird  signs 
and  strange  names  used  by  the  alchemists.  However,  these  attempts 
failed  —  largely  because  they  were  unwieldy  and  inconvenient. 



Berzelius  helps  bring  order  out  of  chaos.  Jons  Berzelius  (ber- 
ze'li-ws)  was  a  Swedish  orphan  who  became,  in  the  words  of  Sir 
Humphry  Davy,  "one  of  the  great  ornaments  of  his  age."  Berze- 
lius invented  a  simple  system  of  chemical  notation  which  he  intro- 
duced in  1814.  Today  it  is  used  in  every  country  of  the  world. 

Berzelius  said  very  logically:  "It  is  easier  to  write  an  abbreviated 
word  than  to  draw  a  figure.  The  chemical  signs  ought  to  be  letters 
for  the  greater  ease  of  writing  and  not  to  disfigure  a  printed  book. 
I  shall  therefore  take  for  the  chemical  sign  the  initial  letter  of  the 
Latin  name  of  each  chemical  element,  thus,  C,  H,  N,  O,  S,  and  P. 
If  the  first  letter  be  common  to  two  metals,  I  shall  use  both  the  initial 
letter  and  another  letter  they  have  not  in  common,  as  gold  (aurum) , 
Au;  silver  (argentum) ,  Ag;  antimony  (stibium) ,  Sb;  tin  (stannum) , 
Sn."  The  first  letter  of  the  symbol  is  always  capitalized;  the  second, 
if  there  is  a  second,  is  not. 

How  a  compound  is  represented  by  a  formula.  When  we  are  deal- 
ing with  a  compound  rather  than  an  element,  we  write  its  abbrevi- 
ated form,  or  formula,  by  placing  side  by  side  the  symbols  of  the 
elements  that  compose  the  compound.  For  example,  the  formula  of 
zinc  oxide,  a  compound  of  zinc  and  oxygen,  is  ZnO;  the  formula  of 
hydrochloric  acid,  a  compound  of  hydrogen  and  chlorine,  is  HC1. 
A  formula  not  only  represents  the  name  of  a  compound,  but  also 
one  molecule  of  that  compound.  Similarly  a  symbol  represents  both 
the  name  of  an  element,  and  one  atom  of  that  element. 

The  use  of  the  subscript.  Berzelius  used  a  subscript,  a  small  num- 
ber placed  below  and  to  the  right  of  a  symbol,  to  indicate  the  number 
of  atoms  of  the  element  represented  by  that  symbol.  Thus,  H2O 
represents  one  molecule  of  water  containing  two  atoms  of  hydrogen 
and  one  atom  of  oxygen.  The  subscript  1  is  never  written. 

As  you  use  Berzelius'  system,  you  will  see  its  remarkable  simplicity 
and  great  value.  However,  it  was  not  accepted  without  a  struggle. 
Even  Dalton  protested  against  it  saying,  "Berzelius'  symbols  are 
horrifying.  A  young  student  might  as  soon  learn  Hebrew  as  make 
himself  acquainted  with  them."  Evidently,  Dalton  must  have  for- 
gotten his  own  very  complex  pictures  of  various  compounds. 

How  is  the  proper  subscript  determined?  You  may  have  already 
wondered  how  it  is  possible  to  tell  what  subscript  to  use.  Why  A12O8 
and  ZnCl2  and  not  A1O  and  ZnCl?  Must  we  memorize  all  formulas  or 
are  there  definite  rules  to  guide  us? 

Fortunately,  it  is  possible  to  write  formulas  without  first  mem- 
orizing them.  To  do  so,  however,  we  must  have  a  thorough  under- 
standing of  valence. 



What  is  valence?  The  law  of  definite  proportions  tells  us  that 
atoms  occur  in  compounds  in  fixed  ratios.  The  following  table  shows 
the  ratio  in  which  hydrogen  combines  with  other  elements  to  form 
four  common  compounds. 




HCI  (hydrogen  chloride) 
H2O  (water) 
NH3    (ammonia) 
CH*    (methane) 



1  chlorine  atom 
1  oxygen  atom 
1  nitrogen  atom 
1  carbon  atom 

In  this  table,  note  the  difference  in  the  number  of  hydrogen  atoms 
with  which  one  atom  of  the  other  elements  combines.  The  valence 
of  an  element  is  the  number  that  tells  us  how  many  atoms  of  hydro- 
gen normally  combine  with  one  atom  of  that  element.  This  is  a 
simplified  definition,  but  will  serve  us  for  the  present. 

From  the  table  we  see  that  one  atom  of  chlorine  combines  with 
one  atom  of  hydrogen  to  form  HCI.  Thus,  we  say  that  chlorine  has  a 
valence  of  one,  or  as  chemists  put  it,  is  monovalent.  One  atom  of 
oxygen  combines  with  two  atoms  of  hydrogen.  So  we  say  that  oxygen 
has  a  valence  of  two,  or  is  divalent.  Nitrogen  has  a  valence  of  three, 
or  is  trivalent.  Carbon  has  a  valence  of  four,  or  is  tetravalent. 

The  idea  of  valence  was  introduced  in  1852  by  Edward  Frank- 
land,  an  English  chemist.  Hydrogen  is  used  as  the  standard,  because 
its  atom  never  combines  with  more  than  one  atom  of  any  other  ele- 
ment. Hence,  if  its  valence  is  considered  to  be  one,  the  valence  of 
every  other  element  must  be  a  whole  number. 

Learning  valences.  There  is  no  "royal  road"  to  the  study  of  va- 
lence. You  will  find  that  memorizing  the  valences  of  the  more  com- 
mon elements  will  save  you  a  tremendous  amount  of  work  and  will 
give  you  a  better  understanding  of  the  material  to  follow.  Table  2 
lists  the  valences  which  are  considered  essential. 

Notice  th'at  in  the  table,  the  symbols  of  the  metals  are  followed 
by  plus  (+)  signs  and  the  nonmetals  and  radicals  followed  by 
minus  signs  (•— ) .  These  signs  represent  electric  charges,  for  reasons 
which  you  will  learn  in  Chapter  16.  The  number  of  these  electric 
charges  correspond  to  the  valences  of  the  element  or  radical.  Group- 



ing  elements  as  metals  and  nonmetals  is  not  a  completely  satisfactory 
means  of  classification  because  some  elements  behave  at  times  as 
metals  and  at  other  times  as  nonmetals. 

Binary  compounds.  A  binary  compound  is  composed  either  of  two 
elements,  two  radicals,  or  one  element  and  one  radical.  If  you  keep 
in  mind  the  following  rules,  you  will  find  writing  the  formulas  for 
binary  compounds  is  a  simple  process.  Practice  is  essential. 










*Ammonium             NH4  * 
Copper  (cuprous)    Cu  + 
Lithium                   Li  * 
Mercury  (ous)         Hg  * 
Potassium               K 
Silver                     Ag  + 
Sodium                   Na  + 

Barium                  Ba  "*"*" 
Calcium                Ca++ 
Copper  (cupric)     Cu  ** 
Iron  (ferrous)         Fe  ** 
Magnesium           Mg"1"1" 
Mercury  (ic)           Hg*^ 
Zinc                      In** 

Aluminum           Al  *++ 
Antimony           Sb"1"*"*1 
Arsenic              As  *  <f+ 
Chromium          Cr  "*"*"*' 
Iron  (ferric)         Fe  "*"*"*" 


Bromine  (bromide)    Br~ 
Chlorine  (chloride)    Cl~ 
Fluorine  (fluoride)     F  " 
Iodine  (iodide)          l~ 

Oxygen  (oxide)    O 
Sulfur  (sulfide)      S~" 

Nitrogen            N 
Phosphorus        P 


*Acetate          C2H3O2~ 
Bicarbonate    HCOs" 
Chlorate         CI03" 
Hydroxide      OH~ 
Nitrate            NOa  " 
Nitrite             NO2  - 

Carbonate            COs 
Sulfate                 S04~~ 
Sulfite                   SO  3  — 

Phosphate          PO4 

*  Radicals:  A  radical  is  a  group  of  atoms  acting  as  a  single  atom  and  having  its  awn  individual 
valence.  The  ammonium  radical,  NH^  has  a  valence  of  ?. 



1 )  Write  the  symbol  with  a  positive  valence  first,  followed  by 
the  symbol  with  a  negative  valence.  Add  plus  and  minus  signs 
to  the  upper*  right  of  the  symbols,  that  is,  in  superscripts,  to 
show  the  valence  for  each  symbol. 

2)  //  the  valences  of  the  symbols  are  equal  no  subscripts  are 
added.  This  rule  is  followed  unless  the  subscripts  represent 
the  actual  structure  of  the  compound.  Thus,  the  formula  of 
hydrogen  peroxide  is  H2O2  and  not  HO,  since  a  molecule  of 
hydrogen  peroxide  actually  contains  two  atoms  of  hydrogen 
and  two  atoms  of  oxygen. 


3)  Since  every  compound  is  electrically  neutral,  the  number 
of  positive  charges  must  be  the  same  as  the  number  of  its 
negative  charges.  Therefore,  //  the  valences  are  not  equal  in 
numerical  value,  subscripts  must  be  added  to  equalize  them. 
Add  to  each  symbol  a  subscript  of  the  same  numerical  value 
as  the  valence  of  the  other  symbol.  The  subscript  1  is  never 

4)  A  radical  acts  like  an  element,  that  is,  it  usually  passes 
thiough  a  chemical  reaction  unchanged.  It  should  be  placed 
in  parentheses  only  if  it  is  followed  by  a  subscript  greater  than 

EXAMPLE  A:  Write  the  formula  for  the  compound,  zinc  oxide. 

1)  Zinc  is  written  Zn  and  has  a  valence  ot  plus  two;  oxygen 
is  written  O  and  has  a  valence  of  minus  two.  The  symbol  lor 
zinc  appears  first  in  the  formula  since  zinc  is  a  metal  and  oxy- 
gen is  a  nonmetal.  Indicate  valences  by  plus  and  minus  signs. 


2)  Since  the  valences  are  equal  no  subscripts  are  written, 
and  the  subscript  for  each  symbol  is  understood  to  be  one. 

3)  The  proper  formula  for  zinc  oxide  therefore  is  ZnO. 

EXAMPLE    B:   Write    the    formula   for    the   compound,    cupric 

1)  Copper  is  written  Cu  and  has  a  positive  valence  of  two; 
chlorine  is  written  Cl  and  has  a  negative  valence  of  one.  Write 
the  symbol  for  copper  first  because  it  has  the  positive  valence. 
Indicate  the  valence  of  each  element  by  using  plus  or  minus 
signs.   (Note  that  the  cupric  valence  of  copper  should  be  used. 
See  Table  2.) 


2)  The  subscript  of  each  symbol  must  be  equal  to  the  va- 
lence of  the  other  symbol.  Since  Cu  has  a  valence  of  plus  two, 
give  Cl  the  subscript  two.  Since  Cl  has  a  valence  of  minus  one, 
the  subscript  for  Cu  is  understood  to  be  one.  The  crossed  ar- 
rows show  these  relationships. 

The  negative  and  positive  valences  are  equal  because  there  is 
one  atom  of  copper  with  a  valence  of  plus  two,  and  two  atoms 
of  chlorine  each  with  a  valence  of  minus  one. 

3)  The   proper   formula   for  cupric   chloride   is   therefore 


EXAMPLE  C:  Write  the  formula  for  the  compound,  magnesium 

1  )  Magnesium  (Mg)  has  a  valence  of  plus  two.  The  sulfate 
radical  (SO4)  has  a  valence  of  minus  two.  Therefore,  Mg  ap- 
pears first  in  the  formula.  Indicate  the  valences  by  using  plus 
and  minus  signs. 


2)  The  valences  of  the  two  symbols  are  equal.  Therefore, 
the  subscript  of  each  is  understood  to  be  one,  and  no  sub- 
scripts are  written. 

3)  The  proper  formula  for  magnesium  sulfate  is  MgSO4. 

EXAMPLE  D:  Write  the  formula  for  the  compound,  zinc  nitrate. 
I  )  Zinc  (Zn)  has  a  valence  of  plus  two.  The  nitrate  radical 
(NO3)  has  a  valence  of  minus  one.  Therefore,  Zn  appears  first 
in  the  formula.  Indicate  the  valences  by  using  plus  and  minus 


2)  Since  Zn  has  a  valence  of  plus  two,  we  give  NO3  the  sub- 
script two  and  enclose  it  in  parentheses.  NO3  has  a  valence  of 
minus  one  and  we  consider  Zn  to  have  the  subscript  one.  The 
crossed  arrows  show  these  relationships. 

3)  The  proper  formula  for  zinc  nitrate  is  Zn(NO3)2. 

Some  elements  have  more  than  one  valence.  Iron  is  divalent 
in  ferrous  compounds  and  trivalent  in  ferric  compounds.  Fer- 
rous chloride  provides  an  example  of  divalent  iron:  Fe++Cl2-. 
Ferric  chloride  provides  an  example  of  trivalent  iron: 

Mercury  is  monovalent  in  mercurous  compounds  and  di- 
valent in  mercuric  compounds;  copper  is  monovalent  in 
cuprous  compounds  and  divalent  in  cupric  compounds.  What 
is  the  meaning  of  -ous  and  -ic  as  related  to  valence? 

In  the  formula  Fe3O4  (Magnetic  oxide  of  iron)  you  might 
think  that  iron  has  a  valence  of  four  and  oxygen  a  valence  of 
three.  However,  the  real  explanation  is  that  this  compound 
is  a  combination  of  Fe++O~-,  in  which  Fe  is  divalent,  and 
Fea+++O8-  ~,  in  which  Fe  is  trivalent.  Oxygen  is  always  divalent. 
Or,  Fe364  may  be  thought  of  as  FeO  •  Fe2O3.  From  these 
formulas  it  is  apparent  that  some  elements  may  exhibit  two  or 
more  valences.  The  electron  theory  offers  an  interesting 


explanation  of  the  fact  that  certain  elements  have  more  than 
one  valence   (see  Chapter  11). 


1.  Write  the  formulas  for  the  following  compounds  showing 
the  +  and  —  signs  and  the  arrows  pointing  toward  the  sub- 
scripts:   (a)  sodium  chloride,    (b)  calcium  bromide,   (c)  ferric 
iodide,   (d)  potassium  fluoride,   (e)  barium  oxide. 

2.  Write  the  formulas  of:   (a)  magnesium  chloride,  (b)  zinc 
oxide,    (c)  aluminum  nitride,    (d)  potassium  sulfide,    (e)  cu- 
prous chloride. 

3.  Write  the  formulas  of:   (a)  aluminum  chloride,  (b)  arse- 
nic oxide,  (c)  calcium  phosphide. 

4.  Write  the  formulas  of:   (a)  sodium  hydroxide,  (b)  potas- 
sium sulfate,  (c)  mercurous  phosphate. 

5.  Write  the  formulas  of:   (a)  calcium  bicarbonate,  (b)    cu- 
prous carbonate,   (c)  magnesium  phosphate. 

6.  Write  the  formulas  of:    (a)  aluminum  nitrate,   (b)  ferric 
sulfate,  (c)  chromium  phosphate. 

7.  Carbon    (C++++)    and  silicon    (Si++++)    are  tetravalent. 
Write  the  formulas  of:   (a)  carbon  tetrachloride  and   (b)  sili- 
con dioxide. 

8.  Write  the  formulas  of:    (a)  mercuric  nitrate,   (b)  sodium 
nitrate,    (c)  mercuric   chlorate,    (d)  mercuric   hydroxide,   (e) 
mercuric  carbonate,    (f)  mercurous  sulfate,    (g)  calcium  sul- 
fite,  (h)  mercuric  phosphate,  (i)  mercurous  chloride. 

How  to  determine  valence  in  compounds  of  more  than  two  ele- 
ments. Remembering  that  in  every  compound  the  number  of  posi- 
tive charges  must  equal  the  number  of  negative  charges,  let  us  find 
the  valence  of  Cr  in  K2CrO4.  There  are  four  oxygen  atoms  each  with 
two  negative  charges  making  a  total  of  eight  negative  charges.  To  bal- 
ance these,  the  compound  has  two  potassium  atoms  each  having  one 
positive  charge  or  a  total  of  two  positive  charges.  The  compound 
must  therefore  have  six  more  positive  charges  which  must  come  from 
the  metal  chromium.  Hence  the  valence  of  Cr  in  this  compound  is 
plus  six. 

Chemistry  has  a  language  and  nomenclature  of  its  own.  Lavoisier 
realized  the  importance  of  language  to  a  science.  In  1789,  the  year 
in  which  the  Bastille  was  stormed,  he  published  a  book  written  in 
the  new  language  of  chemistry.  It  did  not  contain  the  obscure  words, 
the  mystic  symbols,  and  the  pompous  phrases  of  alchemy. 

An  outstanding  nuclear  scientist  is  Dr.  Glenn  T. 
Seaborg  of  the  Radiation  Laboratory,  University  of 
California,  Berkeley,  California.  He  has  played  an 
important  role  in  the  discovery  of  the  trans-uranium 
elements,  numbers  93-100.  What  names  have 
been  given  elements  93-98?  Can  you  suggest 
their  derivation? 

University  <>!  Call  In 

In  naming  the  elements,  several  methods  were  used.  Some,  in- 
cluding bromine,  meaning  stench,  were  named  after  a  physical  prop- 
erty. Some,  including  argon,  meaning  idle,  were  named  after  a  chem- 
ical property.  Some,  including  /;o/onium,  germanium,  gallium,  and 
americium,  were  named  after  countries  or  other  geographic  regions. 
Some  were  named  after  the  city  or  state  connected  with  their  dis- 
covery. Thus  hafnium  was  christened  after  the  Latin  name  for  the 
city  of  Copenhagen,  where  it  was  discovered.  Radioactive  curium  was 
named  after  the  Curies,  who  were  the  earliest  investigators  in  the  field 
of  radioactivity.  77«0rium  and  tantalum  were  named  after  figures 
in  mythology.  The  origins  of  the  names  of  the  earliest  known  ele- 
ments have  been  lost  in  the  darkness  of  antiquity. 

Names  of  metals  and  metallic  radicals  usually  end  in  either  -turn 
or  -um,  as  sodium,  potassium,  platinum,  curium,  hafnium,  alumi- 
num, calcium,  and  the  ammonium  radical. 

Names  of  nonmetals  and  nonmetallic  radicals  usually  end  in  -ine 
or  -gen,  as  chlorine,  bromine,  iodine,  oxygen,  and  nitrogen. 

The  most  common  suffixes  used  in  naming  compounds  are  -ide, 
-ate,  -ite,  -ous,  and  -ic.  The  suflix  -ide  represents  a  binary  compound. 

The  suffixes  -ate  and  -ite  indicate,  as  a  rule,  compounds  of  three 
elements,  one  of  which  is  oxygen.  The  -ite  compound  contains  fewer 
oxygen  atoms  than  the  corresponding  -ate  compound.  Thus  sodium 
sulfate,  Na.,SO4,  contains  four  atoms  of  oxygen,  and  sodium  sulfite, 
Na.,SO3,  three  atoms  of  oxygen.  The  -ate  compounds  are  salts  (see 
page  197)  of  -ic  acids,  and  -ite  compounds  are  salts  of  -ous  acids. 

The  suffixes  -ous  and  -ic  indicate  compounds  in  which  the  metal 
has  a  lower  valence  in  the  case  of  -ous  and  a  higher  valence  in  the  case 
of  -ic.  An  -ous  acid  contains  fewer  oxygen  atoms  than  an  -ic  acid; 
thus  sulfurous  acid  is  H2SO3  and  sulfuric  acid  is  H2SO4;  chlorous 
acid  is  HC1O.,  and  chloric  acid  is  HC1O3. 



Some  commonly  used  prefixes  in  chemistry  are  mono-  (or  uni-) , 
di-  (or  hi-) ,  tri-,  tetra-,  and  pent-.  Mono-  (or  uni-) ,  di-  (or  hi-) ,  tri-, 
tetra-,  and  pent-  stand  for  one,  two,  three,  four,  and  five  atoms.  Thus, 
CO  is  carbon  monoxide,  and  CO2  is  carbon  dioxide.  P2O3  is  phos- 
phorus trioxide  and  P2O5  is  phosphorus  pentoxide.  The  prefix  per- 
means  more  atoms  of  an  element  than  are  found  in  a  more  common 
compound,  and  the  prefix  hypo-  means  less  atoms.  Thus  chloric  acid 
is  HC1O3,  perchloric  acid  is  HC1O4;  chlorous  acid  is  HC1O2,  and 
hypochlorous  acid  is  HC1O. 

In  organic  chemistry,  a  branch  of  chemistry  which  deals  with  the 
more  than  650,000  compounds  of  carbon,  a  more  comprehensive 
nomenclature  has  been  carefully  worked  out.  We  shall  learn  more 
about  organic  chemistry  later. 


Caven,  Robert  M.  and  Cranston,  John  A.  Symbols  and 
Formulae  in  Chemistry,  pp.  1-29.  Blackie  &:  Son,  London, 
1928.  Development  and  use  of  symbols  and  formulas  are  traced 
with  great  clearness  and  interest  in  this  valuable  work. 

Jaffe,  Bernard.  Crucibles:  The  Story  of  Chemistry,  pp.  136- 
156.  Simon  and  Schuster,  New  Ybrk,  1948.  "A  Swede  Tears  Up 
a  Picture  Book"  deals  with  Berzelius'  life  and  his  contributions 
to  the  development  of  chemistry. 

Kendall,  James.  At  Home  among  the  Atoms.  D.  Appleton- 
Century  Co.,  New  York,  1932.  This  eminent  chemist  intro- 
duces in  novel  form  the  problem  of  valences.  He  calls  the 
chapter  "Valencia." 

Oesper,  Ralph  E.  "The  Birth  of  Modern  Chemical  Nomen- 
clature." Journal  of  Chemical  Education,  June,  1945,  pp.  290- 
292.  Aii  interesting  story. 


1.  Each  element  has  a  symbol  that  stands  for  one  atom  of 
the  element. 

2.  Each  chemical  compound  may  be  written  in  abbreviated 
form  as  a  formula  by  placing  the  symbols  of  the  elements  that 
compose  the  compound  side  by  side. 

3.  A  radical  is  a  group  of  elements  that  act  as  a  single  ele- 

4.  Subscripts  ate  used  in  a  chemical  formula  to  indicate  the 
number  of  atoms  of  any  element  which  occur  in  a  molecule  of 


any  substance.  Subscripts  are  also  used  following  the  paren- 
theses around  radical  groups  to  indicate  how  many  radicals 
are  present  in  one  molecule. 

5.  The  valence  of  an  element  is  a  number  that  represents 
the  number  of  atoms  of  hydrogen  with  which  one  atom  of  that 
element  normally  combines  in  forming  a  compound. 

6.  Names  of  metals  and  metallic  radicals  usually  end  in  -turn 
or  -urn;  names  of  nonmetals  often  end  in  -ine  or  -gen. 


Group  A 

1.  What  is  a  symbol? 

2.  Who  introduced  the  modern  symbols  and  formulas  of 

3.  Why  was  it  necessary  to  replace   the  old  alchemical 
symbols  with  a  new  system  of  symbols? 

4.  Write  the  symbols  of  three  gaseous,  one  liquid,  and 
three  solid  elements. 

5.  Several  elements  have  the  same  initial  letter.  How  do 
we  indicate  these  elements  by  symbols? 

6.  What  does  a  formula  indicate? 

7.  (a)  Define  valence,    (b)  Who  introduced  valence  into 

8.  What  element  having  a  valence  of  one  is  used  in  de- 
termining the  valence  of  other  elements? 

9.  What  kind  of  element  would  have  a  valence  of  zero? 
10.  Of  what  importance  to  you  is  a  knowledge  of  valence? 

1  1.  What  is  a  radical? 

12.  Copy  and  complete  the  following  statements:   Metals 
have   .  .  .   valences;  nonmetals  have   .  .  .   valences.  The  most 
common  radical  with  a  plus  valence  is  .... 

13.  Copy  and  complete  the  following  statements:   Three 
metals  with  a  valence  of  one  are   .  .  .  ,   .  .  .  ,  and   ....  Three 
divalent  metals  are  .  .  .,  .  .  .,  and  ....  Three  monovalent  non- 
metals  are  .  .  .  ,  .  .  .  ,  and  ....  The  bicarbonate  radical  has  a 
valence  of  ....  A  radical  with  a  valence  of  three  is  .... 

14.  State  four  rules  for  writing  formulas.  Illustrate  each. 



15.  Give  two  examples  of  elements  that  have  more  than 
one  valence. 

16.  Copy  and  complete  the  following  table.  Do  not  write 
in  this  book. 

Bromide    Sulfide    Chlorate    Sulfate  Phosphate    Oxide  Hydroxide 







17.  Make  a  list  of  the  formulas  of  the  phosphates  of  eight 
different  metals. 

18.  Make  a  list  of  the  formulas  of  the  carbonates  of  eight 
different  metals. 

19.  Make  a  list  of  the  formulas  of  five  ammonium  com- 

20.  Correct  the  following:  FeCl,  CuS2,  Ag(NO3)2,  KSO4, 
Na2C103,  and  NH4  (OH) . 

21.  Write  the  names  and  formulas  of  five  compounds  of 

It    ... 

22.  Name  five  elements  and  tell  how  their  names  were  de- 

23.  How  can  you  tell  whether  the  element  ruthenium  is  a 
metal  or  a  nonmetal? 

24.  Give  the  formulas  and  names  of  two  acids  that  illustrate 
the  difference  in  the  use  of  the  suffixes  ~ous  and  -ic. 

25.  Give  the  formulas  and  names  of  two  compounds  other 
than  acids  that  illustrate  the  difference  in  the  use  of  -ous 
and  -ic. 

I    .  . 

26.  How  would  you  name  the  two  compounds,  BaO  and 


27.  Explain  the  meaning  of  each  letter  and  subscript  in 
these  formulas:  HNO8,  FePO4,  Na2CO8. 

28.  Mark  the  valences  of  the  elements  and  radicals  in  these 
compounds,  using  -{-  and  —  signs.  CuSO4,  HgCl2,  NaClO3, 
Ca  (HC03)  2,  X3  (P04)  2,  (NH4)  8Y,  MnO2. 


29.  Determine  the  valence  of:    (a)  sulfur  in  H2SO4,   (b) 
manganese  in  KMnO4,  and  (c)  chromium  in  K2Cr2O7. 

Group  B 

30.  An  unknown  element  X  has  a  valence  of  three.  Write 
the  formula  for  its  oxide. 

31.  The  metal  Y  has  a  valence  of  two  and  the  nonmetal  Z 
has  a  valence  of  two.  Write  the  formula  for  their  compound. 

32.  The  nomenclature  of  chemistry  is  still  not  completely 
organized.  Can  you  give  any  reasons  for  this  state  of  affairs? 

33.  What   element   was   named   after   a   Finnish   chemist? 
(Read  the  chapter  on  the  rare-earth  elements  in  Weeks'  Dis- 
covery of  the  Elements.) 

34.  Write  a  two-  or  three-page  report  on  the  life  of  Berzelius. 
See  list  of  additional  reading  material. 


1.  Make  a  chart  and,  with  the  help  of  your  teacher  of  art, 
include  as  many  alchemical  symbols  as  you   can  and   their 
modern  equivalents. 

2.  The  old  alchemists  wanted  to  keep  their  knowledge  hid- 
den from  the  rest  of  the  world,  and  so  used  strange  symbols 
and  mysterious  language,   (a)  Name  two  groups  of  people  who 
are  modern  equivalents  of  the  alchemists,   (b)  Tell  how  they 
keep  their  knowledge  to  themselves  and  clothe  their  activities 
with  a  veil  of  mystery,    (c)  What  are  the  reasons  for  their 

3.  Form  a  committee  to  make  a  report  on  the  trades,  pro- 
fessions, and  businesses  in  your  community  that  have  "lan- 
guages of  their  own."    (a)  What  reasons  have  they  for  these 
languages?  (b)  What,  after  all,  is  the  function  of  language? 


THE    OCEAN    OF    AIR 

.  .  .  It  ought  to  be  esteemed  much 
less  disgraceful  to  quit  an  error  for 
a  truth  than  to  be  guilty  of  the  van- 
ity and  perverseness  of  believing  a 
thing  still,  because  we  once  believed 
it.  Robert  Boyle,  1627-1691 

Exploring  the  atmosphere.  Every  day  many  of  us  see  gleaming  air- 
planes flying  swiftly  overhead.  What  are  conditions  up  in  the  air 
where  they  travel?  Of  what  is  the  air  in  which  they  travel  composed? 

Have  you  ever  seen  the  glowing  trail  of  a  meteor  as  it  flashed 
across  the  sky?  Why  did  it  burn?  Have  you  ever  examined  the  fused, 
fire-scarred  surface  of  a  meteorite?  Was  it  hot  when  it  entered  our 
atmosphere,  or  did  it  become  heated  as  it  traveled  through  it? 

Probably  you  know  the  answers  to  these  questions,  but  primitive 
men  would  not  have  been  able  to  answer  them.  The  most  learned  of 
the  alchemists,  even  as  late  as  the  seventeenth  century,  did  not  have 
the  answers.  Some  of  them  believed  that  air  was  empty  space  without 
weight  and  without  substance,  while  others  believed  air  to  be  one  of 
the  "four  elements." 

It  was  not  until  1643  that  Torricelli  (tor-re-chel'le)  invented  the 
mercury  barometer,  which  shows  that  air  has  weight  and  exerts  pres- 
sure. We  know  today  that  approximately  15  pounds  (14.7  Ib.)  of 
air  rest  upon  every  square  inch  of  the  surface  of  the  earth.  Upon 
every  square  foot  rests  a  column  of  air  weighing  more  than  one  ton! 

The  density  of  the  atmosphere  is  not  the  same  throughout  its 
entire  extent.  It  is  densest  at  the  surface  of  the  earth,  and  the  farther 


the  surface  of  the  earth,  the  lighter  the  air  becomes,  At  five 
above  the  earth,  air  pressure  Is  only  5.5  pounds.  At  this  eleva- 
tion the  engine  of  an  airplane  cannot  get  sufficient  oxygen  to  burn  its 
fuel.  U  is  half-starved.  What  it  needs  to  function  properly  is  more 
oxygen,  The  turbosupercharger  supplies  this  by  foreleg  more  air 
into  the  It  has  been  calculated  that  even  at  a  height  of  2000 

miles  there 'is  still  some  air,  although  above  five  miles  normal  breath- 
ing of  human  is  impossible.  About  90  percent  of  the  weight 
of°thf  is  within  12  miles  of  the  earth, 

With  respect  to  temperature,  the  atmosphere  Is  divided  into  sev- 
eral layers, 'The  lowest  layer,  about  six  miles  deep,  is  called  the 
troposphere.  Within  the  troposphere  the  temperature  falls  about 
:   1°F.  with  each,  rise  of  300  feel.  The  layer  above  this  is  called  the. 
!  stratosphere  where  temperature  from  about  —30°  to  —  65°F. 

I  In  the  next  layer,  which  lies  between  18  and  28  miles  above  the 
earth's  surface,  a  temperature  of  about  65  degrees  below  zero  oi^the 
Fahrenheit  scale  is  almost  constant.  Above  these  layers,  at  a  height 
of  between  50  and  200  miles,  are  several  layers  of  electrified  particles 
(of  which  the  Kennelly-Hcaviside  and  the  Appleton  were  the  earliest 
discovered) ,  which  prevent  most  radio  waves  from  out  into 

space,  and"  reflect  them  back  to  the  earth  on  their  way  around  the 
globe.  This  is  called  the  ionosphere,  Its  from. 

A  of  a  cross-section  of  the  showing  fhit  approximate  relative" 

osition  of  the  various  layers. 


morning  to  night  and  varies  with  the  seasons.  The  temperature  in 
these  ionized  bands  rises  with  elevation. 

Extremely  short  radio  waves  of  sufficient  energy  can  penetrate 
these  electrified  layers,  as  was  proved  in  1946  by  Army  Signal  Corps 
scientists,  who  succeeded  in  sending  radar  impulses  to  the  moon. 
Reflections  from  the  moon  were  received.  The  round  trip,  477,714 
miles,  was  made  in  2.4  seconds. 

What  is  air?  Air  is  made  up  of  a  number  of  different  gases.  During 
the  eighteenth  century  three  of  these  —  nitrogen,  oxygen,  and  car- 
bon dioxide  —  were  isolated,  or  obtained  separately,  from  it.  Indi- 
cations were  noted  that  still  other  gases,  less  easily  obtained  in  a 
pure  state,  might  possibly  be  present  in  small  quantities.  These 
gases  were  actually  found  later  on. 

Is  air  a  mixture  or  a  compound?  How  are  the  gases  of  the  air  held 
together?  Are  they  chemically  united  or  are  they  merely  mixed,  just 
as  sand  might  be  mixed  with  clay?  To  answer  these  questions,  sci- 
entists made  use  of  the  law  of  definite  proportions.  If  air  is  a  com- 
pound, they  reasoned,  then  the  composition  of  air  must  be  constant. 

To  determine  whether  or  not  the  composition  of  air  is  constant, 
Dalton  and  others  collected  and  analyzed  samples  of  air  taken  from 
thousands  of  different  places  —  from  the  tops  of  mountains,  over 
lakes,  in  valleys,  in  sparsely  settled  regions,  and  in  congested  areas. 
Gay-Lussac  (ga-lu-sak')  ascended  over  Paris  in  a  hydrogen-filled 
balloon  to  a  height  of  4  miles  to  get  samples  of  air.  Much  more  re- 
cently, rockets  carrying  self-sealing  bottles  brought  down  air  from 
heights  as  much  as  36  miles. 

Although  these  analyses  showed  that  all  the  samples  of  air  varied 
only  slightly  in  composition,  enough  difference  was  noted  to  indicate 
that  the  composition  of  air  is  not  constant.  Hence,  air  could  not  be 
a  compound.  Air,  they  decided,  is  a  mixture  of  gases. 

Some  other  proofs  that  air  is  a  mixture.  The  conclusion  reached 
by  the  early  investigators  that  air  is  a  mixture  was  strengthened  by 
further  evidence.  For  example,  rain  water  contains  air  that  has  been 
dissolved  from  the  atmosphere.  By  boiling  rain  water,  this  air  can  be 
driven  out  and  collected.  Analysis  shows  that  this  air  contains  almost 
twice  as  much  oxygen  as  common  air. 

If  air  were  a  compound,  dissolving  it  in  water  would  not  cause  any 
change  in  its  composition.  But  because  air  is  a  mixture,  each  sub- 
stance of  which  it  is  composed  dissolves  in  water  in  proportion  to  its 
own  solubility  therein.  Since  oxygen  is  more  soluble  in  water  than 
nitrogen,  the  other  main  component  of  air,  dissolved  air  contains 
more  oxygen  than  does  common  air. 





9lass ,     -, 
cylinder ',. 

rise  of  water 

Fig.  18.  Finding  the  per- 
centage of  oxygen  in  the 
air  by  the  use  of  a  eu- 

iron  filings 

"','V:'^*?V  , 

rise  of  water 

When  scientists  succeeded  in  changing  air  into  a  liquid  by  cooling 
it  to  a  temperature  of  about  —  190°C.,  they  found  that  liquid  air 
does  not  have  a  definite  boiling  point.  Instead  of  all  the  liquid  air 
boiling  at  a  definite  temperature,  as  does  water  or  any  other  pure 
compound,  they  found  that  nitrogen  boiled  off  first.  Other  com- 
ponents of  the  air  boiled  off  at  higher  temperatures.  A  pure  com- 
pound has  a  definite  boiling  point,  but  liquid  air  does  not.  This  is 
further  proof  that  air  is  a  mixture. 

Finding  the  percentage  of  oxygen  in  air.  The  percentage  of  oxy- 
gen in  air  is  found  by  using  a  chemical  which  will  react  with  oxygen 
and  remove  it  from  the  air.  For  example,  we  might  place  a  small 
piece  of  white  phosphorus  on  the  coiled  end  of  a  copper  wire  and 
insert  it  in  a  eudiometer,  or  measuring  tube,  which  is  then  inverted 
in  a  cylinder  of  water  (see  Fig.  18)  . 

The  phosphorus,  which  is  easily  oxidized,  soon  combines  with  the 
oxygen  of  the  air  in  the  eudiometer  according  to  the  equation: 

This  reaction  continues  until  the  last  trace  of  oxygen  has  combined, 
forming  phosphorus  trioxide,  white  feathery  crystals  which  readily 
dissolve  in  the  water  in  the  cylinder.  Phosphorus  pentoxide,  P2O5,  is 
formed  also  according  to  the  equation: 

This  is  a  white  solid  which  readily  combines  with  water  and,  hence, 
is  used  as  a  drying  agent,  especially  for  gases. 

As  the  oxygen  unites  with  the  phosphorus,  a  partial  vacuum  is 
formed  in  the  eudiometer,  and  the  greater  air  pressure  outside  the 

*  Also  written,  respectively,  P4OC  and  P4O10. 



tube  forces  water  up  the  tube.  At  the  end  of  a  few  hours  the  chemical 
action  ceases  and  the  water  stops  rising.  The  volume  of  water  that  has 
risen  in  the  eudiometer  is  equal  to  the  volume  of  oxygen  that  was 
originally  present  in  the  air  in  the  measuring  tube.  Because  this  rise 
in  the  height  of  the  water  is  approximately  one-fifth  the  height  of  the 
tube,  this  rise  indicates  that  about  20  percent  of  air  is  pure  oxygen. 
In  performing  this  experiment,  do  not  touch  the  phosphorus. 

Instead  of  phosphorus,  a  less  active  element  such  as  iron  may  be 
used.  To  perform  this  experiment,  the  inside  of  a  eudiometer  is 
moistened  with  water  and  enough  iron  filings  are  added  to  form 
a  thin  layer  of  iron  on  the  inside  walls  of  the  tube,  as  shown  in 
Fig.  18.  The  tube  is  then  inverted  in  water,  and  left  undisturbed 
overnight.  Upon  examination,  it  is  found  that  most  of  the  iron  has 
changed  to  brown  rust,  an  oxide  of  iron,  and  the  water  has  risen 
about  one-fifth  the  height  of  the  eudiometer. 

As  some  molten  metals 
harden  in  the  mold,  tiny 
holes  are  formed  due  to 
the  presence  of  dissolved 
gases.  These  holes  make 
the  casting  unfit  for  use. 

To  prevent  formation  of 
holes  in  the  casting, 
nitrogen  is  bubbled  through 
the  molten  metal  in  the 
mold.  The  nitrogen  carries 
away  the  undesirable 
gases.  This  process  is 
known  as  n/frogen  degas- 

Carbide  and  Carbon  Chemicals  Company 


Nitrogen  is  the  most  abundant  element  in  the  atmosphere.  Nitro- 
gen is  the  principal  gas  that  remains  after  oxygen  has  been  removed 
from  air.  It  was  first  carefully  studied  in  1772,  by  Daniel  Rutherford, 
uncle  of  Sir  Walter  Scott.  Pure  nitrogen  may  be  prepared  in  the 
laboratory  by  gently  heating  ammonium  nitrite,  NH4NO2. 

NH4NO2  -»  N2  T  +  2H2O 

The  physical  properties  of  nitrogen  resemble  those  of  its  partner, 
oxygen,  quite  closely.  It  is  colorless,  odorless,  and  scarcely  soluble 
in  water  (about  two  liters  of  nitrogen  dissolve  in  100  liters  of  water 
under  standard  conditions) .  It  is  even  more  difficult  to  liquefy  than 
oxygen,  requiring  a  temperature  13°C.  lower.  It  is  slightly  lighter 
than  oxygen.  Dry  air  is  about  78  percent  nitrogen  by  volume. 

Chemical  properties  of  nitrogen.  Since  nitrogen  under  normal 
conditions  is  not  chemically  active,  Lavoisier  named  this  element 
azote,  meaning  devoid  of  life.  However,  because  it  is  one  of  the  ele- 
ments found  in  niter  (potassium  nitrate) ,  this  name  was  later 
changed  to  nitrogen. 

Chemically,  nitrogen  differs  completely  from  oxygen.  It  does  not 
burn  under  normal  conditions,  and  does  not  support  respiration. 
It  unites  with  oxygen  only  at  such  temperatures  as  exist  in  the  electric 
arc.  However,  at  higher  and  carefully  controlled  temperatures  and 
pressures,  nitrogen  combines  with  oxygen,  hydrogen,  and  carbon. 
With  calcium,  magnesium,  lithium,  and  wolfram,  it  forms  a  group 
of  compounds  called  nitrides. 

N2  +  O2      -»  2NO  (nitric  oxide) 

N2  +  3H2  — »  2NH3  (ammonia) 

N2  +  2C     — >  C2N2  (cyanogen  —  a  colorless,  poisonous  gas) 

N2  4-  3Mg  — •>  Mg3N2  (magnesium  nitride) 

Life  as  we  know  it  consists  of  plant  and  animal  forms  that  have 
survived  because  they  were  well  suited  to  living  in  the  earth's  at- 
mosphere. Life  on  earth  might  be  radically  different  if  the  oxygen  of 
the  air  were  not  diluted  with  inactive  nitrogen.  Nitrogen  tones  down 
the  chemical  activity  of  the  oxygen  of  the  air.  Hence  burning  and 
other  oxidations  are  not  as  rapid  as  they  otherwise  would  be. 

How  nitrogen  is  used.  The  vast  storehouse  of  free  nitrogen  in  air 
furnishes  an  almost  limitless  supply  of  this  valuable  element.  At- 
mospheric nitrogen  obtained  from  liquid  air  is  used  in  making  am- 
monia, ammonium  compounds,  and  nitrates.  These  nitrogen  com- 
pounds are  essential  in  the  manufacture  of  explosives  and  in 
fertilizers  (see  pages  268,  270,  468,  and  470).  Nitrogen  thus  plays 
a  dual  role,  aiding  both  in  supporting  life  and  in  destroying  it. 


Rayleigh  is  confronted  with  another  riddle  of  the  atmosphere.  In 

1894,  Rayleigh  (ra'li) ,  an  English  scientist,  found  that  "pure"  nitro- 
gen obtained  from  air  weighed  a  little  more  than  an  equal  volume  of 
nitrogen  obtained  from  pure  nitrogen  compounds.  This  puzzled  him. 
Dalton  had  declared  that  the  weight  of  the  atom  never  changed, 
regardless  of  its  source. 

The  difference  in  weight  that  Rayleigh  noticed  was  very  small,  and 
might  have  been  ignored  as  caused  by  experimental  errors.  But  after 
Rayleigh  had  spent  months  investigating  this  problem,  he  became 
convinced  that  the  presence  of  some  other  element  in  the  air  was 
responsible  for  the  difference  in  the  weights  of  the  samples  of  nitro- 
gen. With  the  persistence  of  a  true  research  scientist  he  finally  tracked 
down  the  cause  of  this  difference. 

The  discovery  of  argon,  an  inert  gas  of  the  atmosphere.  Rayleigh, 
together  with  William  Ramsay,  studied  the  experiments  of  Cav- 
endish and  came  across  the  following  statement:  "I  made  an  experi- 
ment to  determine  whether  the  whole  of  the  nitrogen  of  the  at- 
mosphere could  be  changed  to  nitric  acid.  Having  condensed  as 
much  as  I  could,  only  a  small  bubble  of  air  remained.  So  that  if 
there  is  any  part  of  the  nitrogen  of  our  atmosphere  which  differs 
from  the  rest,  we  may  safely  conclude  that  it  is  not  more  than  Y^th 
part  of  the  whole." 

Here  was  a  clue  to  their  problem.  Small  as  this  quantity  was, 
Cavendish  had  not  treated  it  as  negligible  or  as  an  error  in  his  work. 
Rayleigh  and  Ramsay  therefore  repeated  his  experiments  and  isolated 
a  small  volume  of  this  gas  from  the  nitrogen  of  the  air.  After  sub- 
jecting it  to  every  known  test,  they  finally  identified  a  new  element, 
heavier  than  nitrogen,  which,  because  of  its  chemical  inertness,  they 
named  argon,  meaning  lazy. 

Besides  argon,  minute  quantities  of  five  other  inert  gases  are 
now  known  to  be  present  in  the  atmosphere  —  helium  (sun) ,  neon 
(new) ,  krypton  (hidden) ,  xenon  (strange) ,  and  radon  (from  ra- 
dium) .  The  discovery  and  isolation  of  these  gases  from  the  air  are 
other  amazing  examples  of  precise  and  painstaking  research.  When 
we  consider  that  each  of  these  gases  is  colorless,  odorless,  insoluble 
in  water,  and  chemically  inert,  refusing  to  unite  with  even  the  most 
active  elements,  we  begin  to  realize  why  they  eluded  chemists  so  long. 

The  inert  gases  go  to  work.  For  many  years  the  inert  gases  re- 
mained chemical  curiosities.  Helium,  first  identified  in  the  sun  in 
1868  and  later  found  in  considerable  quantities  in  natural  gas,  was 
the  first  to  be  put  to  practical  uses.  It  has  taken  the  place  of  hydtogen 
in  inflating  the  balloons  and  blimps  of  the  Army  and  Navy  and  some 



of  the  weather  observation  balloons  of  the  Weather  Bureau.  When 
mixed  with  oxygen,  it  forms  a  synthetic  air  that  is  used  under  pres- 
sure in  caissons  and  is  supplied  to  deep-sea  divers  to  prevent  the 

The  bends  are  a  type  of  severe  cramps  caused  by  the  sudden  ex- 
pansion and  liberation  of  large  quantities  of  nitrogen  gas  that  have 
entered  the  blood  under  the  great  pressures  to  which  deep-sea  divers 
are  subjected.  When  a  diver  rises  to  the  surface  where  the  air  pres- 
sure is  considerably  lower,  less  helium  is  capable  of  remaining  dis- 
solved in  his  blood,  and,  therefore,  some  of  this  gas  is  liberated  from 
the  blood.  Since  helium  is  40  percent  less  soluble  in  blood  than  is 
nitrogen,  less  helium  gas  will  be  forced  out  of  solution  by  the  decrease 
in  pressure,  and  so  the  substitution  of  helium  for  the  nitrogen  of 
common  air  helps  to  prevent  the  bends. 

Laziness  preferred.  With  the  demand  for  lightweight  metal  parts 
for  airplanes,  brought  on  by  the  tremendous  need  for  aircraft  in 
World  War  II,  research  on  the  problem  of  welding  magnesium  was 
stimulated  and  the  helium-atmosphere  process  was  perfected.  In  an 
atmosphere  of  helium,  oxidation  of  the  magnesium  cannot  take 
place,  and  the  magnesium  or  magnesium  alloys  may  be  welded  with 
ease.  The  development  of  this  method  of  welding  magnesium  is  an 

Linde  Air  Products  Company 

This  workman  is  welding  structural 
aluminum  by  the  sigma  method.  The 
term  tigma  is  derived  from  the  first 
letters  of  the  words  s/tiefaW  inert  gas 
metal  arc. 


example  of  the  way  in  which  the  needs  of  society  and  the  research 
of  scientists  are  related. 

Argon  has  replaced  helium-arc  welding  for  magnesium,  and  other 
metals.  It  is  also  used  to  fill  electric-light  bulbs.  When  an  evacuated 
bulb  is  in  use,  metal  evaporates  from  the  filament,  forming  a  deposit 
on  the  inside  of  the  bulb.  This  deposit  blackens  the  bulb,  making 
it  very  inefficient.  In  an  argon-filled  bulb,  this  evaporation  is  re- 
tarded and  the  lamp  may  be  operated  at  higher  temperatures  than 
if  it  had  been  evacuated  only. 

The  ruby  glow.  The  inert  gases  are  also  widely  used  in  the  glowing 
glass  tubes  so  familiar  in  advertising  signs.  When  an  electric  current 
is  sent  through  a  tube  from  which  air  has  been  removed  and  a  minute 
amount  of  neon  gas  introduced,  the  gas  glows  with  an  orange-red 
light.  The  gas  is  at  low  pressure,  about  12  mm.  of  mercury.  The 
amount  of  current  required  is  extremely  low,  about  ^  of  an  ampere, 
but  the  voltage  varies  between  6000  and  12,000  volts.  Neon  and 
krypton  lights  are  also  used  to  mark  airplane  routes  and  to  signal 
to  airplane  pilots.  Small  neon  glow  lamps  are  used  in  testing  high- 
frequency  electric  circuits  such  as  those  in  radios. 

When  argon  is  used  instead  of  neon,  the  light  produced  is  blue. 
However,  most  of  the  blue  tubes  of  this  type  are  filled  with  mercury 
vapor  rather  than  argon.  Xenon  gives  a  light  blue  light,  and  helium 
a  cream-colored  or  pale  orange  light.  Following  page  382  is  an  il- 
lustration showing  the  colors  produced  by  the  inert  gases  in  lighting. 

The  "idle"  gases  have  thus  been  set  to  work.  A  leader  in  this  field 
predicted  that,  in  time,  much  of  our  lighting  would  be  done  by 
glowing  gas  in  luminous  tubes  instead  of  by  incandescent  filaments. 
Certainly  it  is  true  that  much  outdoor-  lighting  is  now  produced  in 
this  way,  but  the  prediction  would  not  apply  to  indoor  lighting  in 
which  the  field  seems  definitely  being  taken  over  by  a  newer  develop- 
ment, fluorescent  tubes  (see  page  448)  . 

Radon  gas,  enclosed  in  sealed  tubes,  is  used  in  the  treatment  of 
cancer.  The  exact  function  of  the  inert  gases  in  the  air  is  still  not 

Our  atmosphere  also  contains  water  vapor.  Water  vapor  is  always 
present  in  air  in  varying  amounts.  The  waters  of  the  earth  are  con- 
stantly evaporating.  Plants  give  off  immense  quantities  of  water 
vapor  during  transpiration,  and  animals,  too,  exhale  water  vapor. 

Rain,  dew,  snow,  fog,  and  other  similar  phenomena  are  caused  by 
the  condensation  of  the  invisible  water  vapor  of  the  air.  Frequently, 
a  pitcher  of  ice  water  sweats  on  the  outside.  This  sweat  is  the  water 
that  was  formed  when  the  water  vapor  of  the  air  came  in  contact 


with  the  cold  outer  surface  of  the  pitcher  and  condensed.  To  de- 
termine accurately  the  amount  of  water  vapor  in  the  air,  we  can 
pass  a  known  volume  of  air  through  a  drying  agent,  such  as  calcium 
chloride,  CaCl2,  or  phosphorus  pentoxide,  P2O5.  The  increase  in 
weight  in  the  drying  agent  equals  the  weight  of  the  water  vapor  in 
the  sample. 

CaCl2  +  2H2O  ->  CaCl2  2H2O 

Carbon  dioxide,  too,  is  present  in  the  air.  Millions  of  tons  of 
carbon  dioxide  are  poured  into  the  air  daily  by  the  burning  of 
organic  substances,  by  the  decay  of  dead  plant  and  animal  matter, 
and  by  the  breathing  of  living  things. 

C  +  O2-»CO2 

Carbon  dioxide  is  a  colorless,  odorless  gas  fairly  soluble  in  water, 
and  1|  times  as  heavy  as  air.  Because  carbon  dioxide  is  already 
completely  oxidized,  it  does  not  burn.  When  passed  through  lime- 
water  (a  water  solution  of  calcium  hydroxide) ,  it  forms  a  white 
precipitate,  calcium  carbonate,  CaCO3.  This  formation  of  a  white  pre- 
cipitate is  the  common  test  for  carbon  dioxide. 

Ca(OH)2  +  CO2  -»  CaCO3  j  +  H2O 

To  determine  accurately  the  amount  of  carbon  dioxide  present  in 
air,  a  known  volume  of  air  is  passed  through  a  concentrated  solution 
of  potassium  hydroxide,  KOH,  and  the  amount  of  potassium  car- 
bonate, K2CO3,  that  forms  is  determined. 

2KOH  +  CO2  -»  K2CO3  +  H2O 

Carbon  dioxide  and  its  uses  are  discussed  in  Chapter  23. 

Cause  and  effect  of  atmospheric  pollution.  An  adult  inhales  about 
37  pounds  of  air  a  day,  which  is  five  times  the  weight  of  the  food  and 
water  that  he  consumes.  We  are  very  careful  about  getting  pure 
water,  and  have  laws  to  protect  us  against  the  sale  of  impure  food. 
But  we  have  not  done  as  much  about  polluted  air.  In  some  cities 
legislation  has  been  passed  to  cut  down  pollution  of  the  air  caused 
by  smoke.  Excessive  smoke  may  impair  health,  damage  crops,  slowly 
destroy  property,  and  reduce  visibility.  Smog,  a  combination  of 
smoke  and  fog,  is  another  serious  problem  (see  pages  329,  616,  and 

What  is  air  conditioning?  Not  so  many  years  ago,  it  was  supposed 
that  the  air  in  a  crowded  room  was  unhealthful  because  it  contained 


a  large  percentage  of  carbon  dioxide  and  a  lowered  percentage  of 
oxygen.  It  has  since  been  proved  that  even  in  a  very  crowded  room 
the  percentage  of  carbon  dioxide  never  reaches  a  point  where  it  be- 
comes harmful.  The  amount  of  oxygen  rarely  gets  below  20  percent, 
and  it  can  be  cut  down  even  to  17  percent,  at  which  point  a  candle 
is  extinguished,  without  being  injurious  to  health. 

Research  has  shown  that  "bad  air"  is  really  caused  by  high  temper- 
ature, lack  of  circulation,  high  percentage  of  water  vapor,  and  various 
odors  that  have  accumulated. 

Warm  air  can  hold  more  water  vapor  than  an  equal  volume  of  cold 
air.  A  cubic  meter  of  air  at  20°C.  (68°F.) ,  for  example,  is  capable 
of  holding  about  17  grams  of  water  vapor,  while  the  same  volume  of 
air  at  11°C.  (52°F.)  can  hold  only  about  one-half  as  much.  The 
ratio  of  the  weight  of  water  vapor  present  in  air  to  the  weight  of 
water  vapor  it  is  capable  of  holding  under  the  existing  conditions  of 
temperature  and  pressure  is  known'  as  its  relative  humidity. 

The  temperature  of  the  human  body  is  controlled  in  part  by  the 
evaporation  of  perspiration.  Evaporation  absorbs  heat.  The  amount 
of  heat  absorbed  depends  upon  the  amount  of  perspiration  evap- 
orated. The  cooling  sensation  produced  depends  upon  both  the 
amount  and  the  rate  of  evaporation. 

Air  with  a  high  relative  humidity,  regardless  of  temperature,  will 
evaporate  little  perspiration.  Consequently,  on  hot,  damp  days  we 
feel  hotter  than  we  do  on  hot  days  ihat  are  somewhat  drier. 

Tests  have  shown  that  in  winter  the  air  in  many  homes  has  a  very 
low  relative  humidity.  Such  very  dry  air  has  a  twofold  bad  effect. 
It  tends  to  "dry  us  out."  Evaporation  of  perspiration  present  proceeds 
rapidly  with  great  cooling  effect.  Consequently,  the  temperature  of 
dry  air  must  be  greater  than  the  temperature  of  damper  air  for  our 
sensations  of  warmth  to  be  the  same.  It  t^kes  more  fuel  to  maintain 
the  higher  temperature  and,  as  a  result,  costs  more.  In  addition, 
temperatures  above  21°C.  (70°F.)  are  likely  to  produce  drowsiness 
and  prevent  us  from  doing  our  best  work.  Furthermore,  such  dry  air 
tends  to  dry  out  the  linings  of  the  nose,  mouth,  and  throat,  thus 
causing  great  discomfort  and  reducing  our  resistance  to  common 
colds  and  other  respiratory  diseases. 

Today,  steady  progress  is  being  made  in  supplying  properly  con- 
ditioned air,  not  only  to  large  auditoriums,  factories,  classrooms, 
banks,  office  buildings,  and  railroad  trains,  but  even  to  subway  cars 
and  private  homes.  Air-conditioning  equipment  filters  out  dust  and 
pollen  from  the  air,  exhausts  stale  air,  keeps  the  relative  humidity 
at  the  right  point  (about  50  percent) ,  and  maintains  a  comfortable 



TABLE  3.  I 




temperature  (about  68 °F.) .  It  also  destroys  the  dulling  quality  of 
"dead"  air  by  keeping  the  air  in  motion.  In  addition  to  increasing 
mental  and  physical  efficiency,  air  conditioning  is  essential  in  certain 
manufacturing  processes,  such  as  printing  and  the  making  of  pills, 
chocolate,  rayon,  paper,  tobacco  products,  and  steel  bearings. 

Summary:  The  composition  of  the  atmosphere.  The  chief  con- 
stituents of  samples  of  dry  air  taken  near  the  surface  of  the  earth  are 
shown  in  the  following  table. 





,                 Oxygen 
Carbon  dioxide 



Air  contains  variable  quantities  of  water  vapor  and  carbon  dioxide; 
small  amounts  of  the  inert  gases,  argon,  helium,  krypton,  neon, 
radon,  and  xenon;  and  also  minute  amounts  of  other  gases,  such  as 
methane,  carbon  monoxide,  hydrogen,  nitrogen  dioxide,  and  ozone, 
as  well  as  very  finely  divided  solids,  such  as  dust,  bacteria,  spores,  and 

Liquid  air.  In  Gullivers  Travels,  a  famous  Academy  was  visited, 
and  Jonathan  Swift  reports  how  some  of  its  scientists  were  con- 
densing air  and  letting  the  liquid  flow  like  water.  Probably  Swift 
believed  the  liquefaction  of  air  a  dream  never  to  be  realized.  Yet 
today  liquid  air  is  a  common  article  of  commerce  and  thousands  of 
tons  of  it  are  used  every  year. 

The  principles  underling  the  manufacture  of  liquid  air  are: 
(1)  when  a  liquid  evaporates,  it  absorbs  heat  from  its  surroundings 
and  thereby  lowers  their  temperature;  and  (2)  the  sudden  expansion 
of  a  gas  produces  this  same  effect,  When  air  is  compressed,  cooled, 
and  suddenly  allowed  to  expand  through  a  narrow  opehing,  its 
temperature  is  lowered.  If  this  process  is  repeated,  a  temperature  is 
finally  reached  which  is  low  enough  to  liquefy  the  gas.  A  more  de- 
tailed treatment  of  the  liquefaction  of  gases  will  be  found  on 
page  253. 

How  the  kinetic  theory  of  matter  explains  the  liquefaction  of 
air.  All  matter  is  thought  to  be  made  up  of  small  particles  (atoms 
and  molecules)  that  are  in  constant  motion.  In  the  case  of  gases,  this 
motion  is  extremely  rapid  (air  molecules  under  normal  conditions 



move  at  about  20  miles  per  minute) .  When  any  gas  is  cooled,  its 
molecules  move  more  slowly,  until  finally  a  temperature  is  reached 
at  which  the  motion  of  the  molecules  of  gas  is  so  slow  that  they  come 
close  enough  together  to  form  larger  groups  or  clusters  of  molecules, 
thus  forming  a  liquid.  When  a  gas  expands,  some  heat  is  used  in  sep- 
arating the  molecules. 

Compressing  a  gas  has  the  effect  of  bringing  the  molecules  closer 
together.  By  a  simultaneous  cooling  and  compressing,  any  gas  may 
be  liquefied.  Some  gases,  such  as  chlorine,  sulfur  dioxide,  and  am- 
monia, are  easily  liquefied.  Other  gases,  such  as  oxygen,  nitrogen, 
hydrogen,  and  helium,  require  extremely  low  temperatures  and 
high  pressures  to  change  them  to  liquids.  It  has  been  calculated  that 
at  — 273°C.  (absolute  zero)  all  motion  of  the  particles  of  matter 
ceases.  The  nearest  approach  to  this  temperature  thus  far  attained 
was  made  in  1952  when  scientists  at  our  National  Bureau  of  Stand- 
ards reached  a  temperature  within  0.001,5°  of  absolute  zero.* 

*  The  measurement  of  absolute  temperatures  is  discussed  on  p.  644. 

Arthur  D.  Little,  Inc. 

Filling  a  flask  with  liquid  helium 
in  an  experimental  laboratory. 
Helium  is  liquefied  by  subjec- 
ting it  to  both  cooling  and  pres- 




19.   A  thermos   bottle 
a   Dewar  flask.    Note 
imilarities.  What  is  the 
of  the  silvered  sur- 

Properties  and  uses  of  liquid  air.  Liquid  air  is  a  pale  blue  liquid 
almost  as  heavy  as  water.  It  contains  about  21  percent  oxygen  and 
boils  at  —  190°C.  When  it  evaporates,  nitrogen  boils  off  first  and  the 
mixture  becomes  richer  and  richer  in  oxygen,  for  the  same  reason 
that  the  boiling  alcohol-water  mixture  in  your  car's  radiator  loses 
alcohol  faster  than  water,  because  the  boiling  point  of  alcohol  is 
lower  than  that  of  water. 

Liquid  air  is  used  chiefly  as  a  source  of  oxygen,  nitrogen,  and  the 
inert  gases,  which  boil  off  at  different  temperatures.  The  Nazis  used 
liquid  oxygen  and  alcohol  to  fuel  the  V-2  rocket  bombs  which  they 
hurled  against  England  in  1945. 

Because  of  the  great  tendency  of  liquid  air  to  evaporate,  small 
volumes  of  it  are  kept  in  special  containers  called  Dewar  flasks,  which 
are  similar  in  construction  to  the  familiar  thermos  bottle.  These 
flasks  cannot  be  tightly  stoppered,  for  any  attempt  to  confine  liquid 
air  too  closely  results  in  explosion.  Because  of  the  danger  of  injury 
both  from  such  explosions  and  from  contact  with  a  substance  at  such 
an  extremely  low  temperature,  persons  handling  liquid  air  must  use 
great  care. 

Air  Reduction  Company,  Inc. 

Inferior  of  an  oxygen-nitrogen 
plant.  The  workman  is  standing  in 
front  of  a  column  in  which  the  two 
gases  are  obtained  from  liquid  air. 


The  properties  of  substances  change  when  immersed  in  liquid  air. 
Liquid  mercury  becomes  solid  enough  to  be  used  as  a  hammer  head. 
Rubber  turns  hard  and  brittle.  The  resistance  of  copper  to  an  elec- 
tric current  is  decreased  50  times. 


Cady,  H.  P.  "Liquid  Air."  Journal  of  Chemical  Education, 
June,  1931,  pp.  1027-1043.  A  fascinating  account  of  experi- 
ments with  liquid  air. 

Kaempffert,  Waldemar.  Explorations  in  Science,  Chapter  7, 
pp.  102-108  is  entitled,  "This  most  excellent  canopy,  the  air." 
The  Viking  Press,  New  York,  1953. 

Ramsay,  William.  The  Gases  of  the  Atmosphere,  pp.  148- 
181,  234-269.  The  Macmillan  Co.,  London,  1915.  Sir  William 
tells  of  the  discovery  of  argon  and  discusses  other  inert  gases. 


1.  Air  is  a  mixture  of  gases  because   (1)   it  has  no  definite 
composition,    (2)    when  air  dissolves  in  water,  the  dissolved 
air  contains  more  oxygen  than  common  air,  and  (3)  liquid  air 
does  not  have  a  definite  boiling  point. 

2.  Air  conditioners  cleanse  the  air,  keep  the  relative  humid- 
ity where  it  belongs   (about  50  percent) ,  maintain  a  comfort- 
able temperature  (68  °F.) ,  and  keep  the  air  in  motion. 

3.  The  kinetic  theory  of  matter  assumes  that  gases  are  made 
up  of  small  particles  (atoms  and  molecules)  in  active  motion. 
Cooling  the  gas  slows  down  this  motion  and  brings  the  par- 
ticles closer  together  until  a  liquid  is  formed.  The  molecules 
of  liquids  and  solids  are  also  in  motion. 


Group  A 

1.  Who  first  definitely  proved  that  air  has  weight? 

2.  (a)  What  pressure  does  the  atmosphere  exert  at  sea 
level?  (b)  What  causes  this  pressure? 

3.  What  are  the  differences  between  the  troposphere,  strat- 
osphere, and  Heaviside  layer? 

4.  Does  the  composition  of  the  atmosphere  prove  that  air 
is  a  mixture?  Explain. 

5.  Give  two  reasons  other  than  composition  for  believing 
that  air  is  not  a  compound. 


6.  State  the  composition  of  dry  air  by  volume. 

7.  What  substances  are  found  in  air  in  variable  quantities? 

8.  Make  a  diagram  illustrating  a  laboratory  method  for 
determining  the  percentage  of  O2  in  air. 

9.  (a)  What  is  the  gas  that  is  left  in  the  measuring  tube 
used  in  determining  the  percentage  of  O0  in  air?    (b)  What 
are  the  impurities  in  this  gas? 

10.  State  four  physical  properties  of  N2. 

11.  What  is  the  chief  chemical  property  of  N2? 

12.  Why  is  it  wrong  to  call  N2  an  inert  gas? 

13.  Write  a  chemical  equation  illustrating  the  action   of 
hot  Mg  on  N2. 

14.  Under  what   conditions  does   N.,   combine   chemically 
with  O3  and  H2?  » 

1  I 

15.  (a)  State  two  functions  of  the  N2  of  air.   (b)  State  two 
commercial  uses  of  N2. 

16.  (a)  Why  is  Mg  difficult  to  weld?   (b)  Under  what  con- 
dition is  it  easily  welded?    (c)  Development  of  this  process 
is  an  example  of  what? 

17.  Name  the  six  inert  gases  of  the  atmosphere. 

18.  What  five  properties  are  common  to  all  six  of  the  inert 
gases  of  the  atmosphere? 

19.  Match  each  element  listed  in  the  first  column  with  the 
correct  item  in  the  second  column. 

1)  Ne  a)   hidden 

2)  A  b)  balloons 

3)  Kr  c)  strange 

4)  He  d)  luminous  tubes 

5)  Xe  e)  electric-light  bulbs 

f)  airplanes 

20.  Explain  the  "sweating"  of  a  pitcher  filled  with  ice  water, 

21.  How  would  you  determine  the  amount  of  water  vapor 
present  in  air? 

22.  How  would  you  determine  the  amount  of  CO,  present 
in  air? 

23.  State  (a)  four  physical  and  (b)  two  chemical  properties 
of  C02. 

24.  What  is  air  conditioning? 

25.  Name    (a)  two  causes  and    (b)  three  effects  of  atmos- 
pheric pollution. 


26.  What  two  conditions  in  the  air  of  a  poorly  ventilated 
room  make  it  both  uncomfortable  and  unhealthful? 

t .  .  . 

27.  What  is  relative  humidity? 

28.  Upon   what  principle  does   the   liquefaction  of  a  gas 

29.  Discuss  the  meaning  of  the  kinetic  theory  of  matter. 

30.  What  is  the  chiet  use  of  liquid  air? 

31.  State  three  properties  of  liquid  air. 

32.  Compare  the  effectiveness  of  determining  the  percentage 
of  O2  in  air  by  the  phosphorus  and  the  iron  filings  methods. 

Group  B 

33.  Why  does  the  composition  of  air  by  weight  differ  from 
its  composition  by  volume? 

34.  A  piece  of  burning  charcoal  is  plunged  into  liquid  air. 
It  keeps  on  burning  with  even  greater  splendor.  Why  does  not 
the  extreme  cold  of  liquid  air  extinguish  the  burning  charcoal? 

35.  Explain  the  presence  of  vast  amounts  of  free  N2  in  air, 
and  only  relatively  small  amounts  of  nitrogen  compounds  in 
the  earth's  crust. 

36.  In  different  samples  of  air  the  following  substances  are 
placed:  CaCl,,  P,  hot  Mg,  and  Ca  (OH) ,.  Explain  what  hap- 
pens in  each  case. 

37.  Why  does  a  kettle  of  liquid  air  boil  when  placed  on  ice? 

38.  "Coal   burned   in   our  furnace   returns   to   us   in   our 
bread."  Explain. 

39.  What  weight  of  dry  air  would  be  theoretically  needed 
to  extract  ten  grams  of  pure  oxygen? 


1.  Visit  a  neon-sign  factory  or  an  air-conditioned  building 
or  factory  and  make  a  full  report  on  your  visit  to  the  class. 

2.  Prepare   simple    demonstrations    to   show    that    (1)  air 
contains  water  vapor,  and  (2)  your  breath  contains  CO2.  Per- 
form these  experiments  before  your  class  the  following  day. 

3.  With  the  aid  of  a  wet  and  dry  bulb  thermometer,  de- 
termine the  relative  humidity  of  your  classroom. 

4.  Prepare  a  report  on  the  probable  origin  of  the  earth's 
atmosphere.  Use  a  good  book  on  physiography  or  one  of  the 
references  listed. 





.  .  .  Here  evidently  we  are  at  the 
birthplace  of  the  chemical  equation, 
yet  we  cannot  find  in  the  writings  of 
Lavoisier  this  instrument  as  we  know 
it;  for  our  chemical  equation  de- 
pends as  much  on  the  atomic  theory 
as  on  the  doctrine  of  the  conservation 
of  mass.  Caven  and  Cranston,  1928 

Chemical  equations,  the  shorthand  of  chemistry.  Chemists  use 
symbols  and  formulas  that  are  understood  to  represent  elements  and 
compounds  by  scientists  in  all  parts  of  the  world.  But  chemists  make 
further  use  of  these  symbols  and  formulas.  They  use  them  to  tell  the 
story  of  chemical  change  —  the  reacting  substances,  the  type  of  chem- 
ical change,  the  products  of  the  reaction,  and  various  other  facts. 
All  of  these  facts  are  expressed  in  the  form  of  a  chemical  equation, 
which,  in  certain  respects,  is  similar  to  the  equations  you  have  used 
in  arithmetic  and  algebra. 

Considering  the  chemical  reaction  that  takes  place  when  iron  and 
sulfur  are  heated  together,  chemists  and  students  of  chemistry  write 
as  follows: 

Fc  +  S  -»  FcS 

As  you  see,  the  two  sides  of  the  equation  are  separated  by  an  arrow, 
-».  This  arrow  means  yields,  or  forms,  and  should  be  read  using 
these  or  similar  words.  On  the  left  side  of  the  arrow,  chemists  write 
the  symbols  or  formulas  of  the  substances  that  react.  On  the  right  side 
of  the  arrow,  they  write  the  symbols  and  formulas  of  the  products  of 
the  reaction.  A  word-equation  that  will  express  the  same  reaction 
as  the  chemical  equation  that  we  have  just  been  discussing  follows: 




Iron  reacts  with  sulfur  to  yield  iron  sulfide. 

Equations  represent  reality.  To  write  a  chemical  equation,  we 
must  know  the  formulas  of  the  substances  involved.  To  write  a 
chemical  equation  correctly,  we  must  know  exactly  what  substances 
are  reacting  and  what  substances  are  the  products  of  the  reaction. 
The  facts  are  determined  before  an  equation  is  written. 

We  can  write  equations  for  only  those  chemical  reactions  known  to 
be  capable  of  actually  taking  place.  Chemical  equations  must  repre- 
sent actual  conditions.  For  example,  it  is  incorrect  to  write  He  +  S 
-»  HeS,  because  helium  and  sulfur  do  not  react  together,  and 
helium  sulfide,  HeS,  has  never  been  prepared. 

Balancing  an  equation.  Lavoisier  once  wrote:  "One  may  take  it 
-for  granted  that  in  every  reaction  there  is  an  equal  quantity  of  matter 
before  and  after  the  operation."  Atoms  do  not  disappear  in  the 
process  of  chemical  reaction.  Therefore,  the  number  of  atoms  in 
the  reacting  substances  must  equal  the  number  of  atoms  in  the 
products  of  the  reaction.  The  number  of  atoms  of  each  element  must 
be  the  same  on  each  side  of  the  arrow.  Every  correct  equation  con- 
forms, in  this  way,  to  the  law  of  the  conservation  of  matter. 

Points  to  remember  in  balancing  equations. 

1)  Elements  can  occur  in  a  free  state,  not  combined  with  other 
elements.  For  reasons  which  you  will  learn  later  (pages  279—280)  , 
a  number  of  these  elements,  including  oxygen,  hydrogen,  nitrogen, 
chlorine,  bromine,  and  iodine  occur  as  molecules  containing  2  atoms. 
They  are  written  O,,  H.,.  N2,  C12,  Br2,  and  I2,  thus: 

H2  +  C12  —  »  2HC1  (hydrogen  chloride) 
N2  4-  O2  —  >  2NO  (nitric  oxide) 

.  .     <  Cu  4-  Br2  —  >  CuBr2  (copper  bromide) 

2H2O  ->  2H2  +  02 

Standard   Oil   Company 

Delicate  measurements  must  be 
made  to  determine  the  actual 
conditions  that  are  described 
by  chemical  equations. 


2)  In  general,  a  radical  remains  unaltered  during  a  chemical 
change  and  its  symbols  are  carried  over  to  the  right  side  of  the  equa- 
tion unchanged: 

Zn  +  H2SO4  ->  ZnSO4  +  H2 
2Na  +  2HOH  -»  2NaOH  +  H2 
BaO2  +  H2SO4  ->  BaSO4  +  H2O2 

Balancing  an  equation  consists  of  changing  coefficients  until  the  same 
total  number  of  atoms  of  each  element  is  shown  on  each  side  of  the 

Do  not  alter  the  subscript  of  a  radical  nor  any  other  subscript  in 
order  to  make  an  equation  balance.  Such  alteration  would  mean  a 
change  in  the  actual  composition  of  the  compound.  This  in  turn 
would  mean  an  entirely  different  substance  which  does  not  actually 
appear  in  the  chemical  change.  We  may,  however,  alter  coefficients 
of  any  element  or  compound  without  changing  the  meaning  of  the 
chemical  symbols.  The  coefficient  describes  the  number  of  molecules 
of  a  substance.  By  changing  the  number  of  molecules  present  in  a 
chemical  change,  we  do  not  alter  the  composition  of  the  substances 
which  are  involved. 

Your  best  preparation  for  writing  equations  correctly  is  a  thorough 
knowledge  of  valence  and  of  formulas.  In  addition,  you  should 
attack  the  problem  in  a  systematic  and  thoughtful  way.  The  follow- 
ing procedure  should  be  helpful. 


1)  Write  the  equation  without  giving  coefficients  to  any  of 
the  formulas. 

2)  Write  any  free  element  occurring  in  the  equation  without 
a  subscript.  Retain  all  other  subscripts. 

3)  Select  the  compound  with  the  greatest  number  of  atoms. 
For  one  of  its  elements,  compare  the  number  of  atoms  appear- 
ing on  each  side  of  the  equation.  If  the  numbers  differ,  decide 
upon  a  coefficient  or  coefficients  which  will  equalize  the  num- 
ber of  atoms  on  each  side. 

4)  Repeat  this  procedure  for  the  other  elements  until  the 
equation  is  balanced,  that  is,  until  the  same  number  of  atoms 
of  each  element  appears  on  each  side  of  the  equation. 

5)  Add  the  subscript  to  any  free  element  in  the  equation 
which  may  require  it.  At  this  stage,  free  elements  may  have 
been  given  coefficients.  If  the  coefficient  of  the  free  element  is 


an  even  number  divide  the  coefficient  by  two  when  the  sub- 
script is  added.  The  total  number  of  atoms  of  the  free  element 
will  then  remain  the  same  (for  example,  6H  contains  the  same 
number  of  atoms  as  3H2) . 

6)  If,  however,  the  coefficient  of  the  free  element  is  an  odd 
number,  it  cannot  be  divided  by  two.  In  this  case,  when  the 
subscript  is  added  to  the  free  element,  all  other  formulas  on 
both  sides  of  the  equation  must  be  multiplied  by  two. 

EXAMPLE  A:  Write  the  equation  for  the  reaction  which  takes 
place  when  potassium  chlorate  is  heated  to  form  potassium 
chloride  and  oxygen. 

1)  Write  the  equation  without  coefficients.  Note  that  the 
free  oxygen  which  is  usually  written  O2  is  written  at  this  stage 
without  a  subscript. 

KClOa  -^  KC1  +  O 

2)  KC1O3   is   the  compound  with   the   largest  number  of 
atoms.  By  inspection,  we  see  that  oxygen  is  the  only  element 
in  the  compound  which  does  not  appear  with  an  equal  number 
of  atoms  on  the  other  side  of  the  arrow.  To  make  the  O  on  the 
right  balance  the  O3  on  the  left,  we  multiply  O  by  the  coeffi- 
cient 3: 

KC1O3  ->  KC1  +  3O 

3)  As  we  have  just  learned,  free  oxygen  must  be  written  as 
•O.,.  Add  the  subscript  to  the  free  oxygen  on  the  right  of  the 
arrow.  The  coefficient  3  is  not  divisible  by  two,  so  when  we 
add  the  subscript,  we  find  there  are  six  atoms  of  free  oxygen 
on  the  right  side  of  the  arrow,  but  only  three  atoms  of  oxygen 
on  the  left  side: 

KC1O3  ->  KC1  +  3O2 

4)  We  may  make  the  equation  balance  again  by  multiplying 
.all  other  formulas  by  two: 

2KC1O3  ->  2KC1  +  3O2 

EXAMPLE  B:  Write  the  equation  for  the  burning  of  benzene 
(C6H0)  to  form  CO2  and  H2O. 

1)  Write  the  unbalanced  equation.  Do  not  write  the  sub- 
.script  for  the  free  oxygen. 

O  ->  CO2  +  H2O 


2)  C6H6  is  the  compound  with  the  largest  number  of  atoms. 
It  contains  6  atoms  of  carbon,  but  there  is  only  one  atom  of 
carbon  to  the  right  of  the  arrow.  Therefore,  we  place  the 
coefficient  6  in  front  of  CO. 

3)  C6H6  also  contains  6  atoms  of  hydrogen,  but  there  are 
only  2  atoms  of  hydrogen  on  the  right  side  of  the  arrow.  There- 
fore, we  place  the  coefficient  3  in  front  of  H2O. 

CeH6  +  O  ->  6CO2  +  3H2O 

4)  We  now  see  that  there  are  15  atoms  (6  X  2  +  3  X  *)  of 
oxygen  on  the  right,  but  only  one  atom  of  oxygen  on  the  left. 
Therefore,  we  place  the  coefficient  15  in  front  of  O. 

C6Hfl  +  150  ->  6C02  +  3H2O 

5)  Since  free  oxygen  is  written  O2,  we  add  the  subscript. 

6)  Since  15  is  not  divisible  by  two,  we  may  bring  the  equa- 
tion into  balance  once  again  by  multiplying  all  other  formulas 
by  the  coefficient  2. 

SQHe  +  15O2  ->  12CO2  +  6H2O 

If  the  coefficient  of  the  oxygen  had  been  an  even  number,  we 
could  have  divided  it  by  two  when  adding  the  subscript.  It 
would  then  be  unnecessary  to  multiply  the  other  formulas  by 
the  coefficient  2. 

HOW  TO  CHECK  EQUATIONS.  To  check  an  equation,  proceed  as  follows: 

1)  Examine  the  first  symbol  on  the  left.  Compare  the  num- 
ber of  atoms  against  the  number  of  atoms  of  the  same  element 
on  the  other  side  of  the  equation.  If  the  numbers  are  equal, 
put  a  small  check  over  the  symbols  for  the  element. 

2)  Proceed  to  each  symbol  in  turn  until  you  have  put  a  check 
over  every  symbol  in  the  equation. 

Thus,  in  checking  the  final  equation  in  Example  A  above, 
2KC1O8  -*  2KC1  +  3O2,  examine  the  first  symbol,  K.  There  are  two 
atoms  of  potassium  on  each  side.  Put  a  check  over  the  K  on  each  side. 
Examine  the  next  symbol,  Cl.  Again  we  have  two  atoms  on  each 
side.  Check  them.  Finally,  we  have  six  atoms  of  oxygen  on  each  side, 
which  we  check.  All  the  symbols  are  then  checked  and  accounted  for: 

v/  v/  v/  V  \/  v/ 

2KC1O3  ->  2KC1  +  3O2 



The  equation  is  now  balanced  accurately.  It  is  correct  both  mathe- 
matically and  chemically.  Following  the  same  procedure,  check  the 
final  equation  in  Example  B. 

Fig.  20.  The  of  th«  tuft  contains  o 

of  oncf  an  test  holding  a 

of  AgN03.  When  the  |t  inverted 

(right),  A§CI  is  but  the  cca|e  tho 

no    change.    No    matter    hot 





To  be  sure  that  you  understand  the  suggestions  just  given  for  bal- 
ancing equations,  balance  the  following: 

a)  NaClOs  ->  NaCl  +  O  g) 

b)  Hg  +  0-»HgO  h) 

c)  P205  +  H20 ->  HPO,  i) 

d)  Mg  +  0->MgO  j) 

e)  Na  +  Cl->NaCl  k) 

f)  C  +  O->CO 

Gu  +  O  ->  CuO 
Fe  -f  O  —  >  Fe2O3 
O  +  P  -»  P20» 
P  4-  O  ->  P2O5 
SiO2  -f  G  -*  SiC 


What  are  the  four  types  of  chemical  reactions?  There  are  four 
general  types  of  chemical  reactions. 

1)  Direct  combination,  or  synthesis.  When  two  or  more  elements 
or  compounds  combine  directly,  forming  a  chemically  more  complex 


t .  . . 

6.  Why  must  subscripts  never  be  changed  in  balancing 

7.  What  is  wrong  with  the  following: 

2Ne   (neon)   +  O2  -»  2NeO? 

8.  Why  may  coefficients  be  changed  in  balancing  an  equa- 

9.  Why  is  a  knowledge  of  valence  so  essential  in  writing 

10.  Name  five  elements  which,  when  free,  must  be  written 
with  the  subscript  2.  ft 

a  1 

11.  Correct  this:  Mg  -(-  Br  -»  MgBr2. 

12.  Generally,  what  happens  to  a  radical  during  a  chemical 

13.  Balance  the  following  equation,  giving  your  reason  for 
each  step: 

Acetylene  (C2H2)   -f  oxygen  (O2) 

->  carbon  dioxide   (CO2)    -\-  water   (H2O) 

14.  What  are  the  four  general  types  of  reactions? 

15.  Give   an   example   of   an   equation   illustrating   direct 

16.  Give  an  example  of  an  equation  representing  simple 

17.  How  do  you  recognize  an  equation  representing  simple 

18.  Give  an  example  of  an  equation  representing  double 

19.  What  is  the  general  rule  followed  in  balancing  a  double- 
replacement  equation? 

20.  Complete  and  balance  the  following  equation.   Show 
how  to  check  it  to  see  if  it  is  correctly  balanced. 

Na  -f  HOH  -»  NaOH  +  H 

21.  Show  in  detail  how  to  balance  the  equation  represent- 
ing the  decomposition  of  sodium  chlorate,  NaClO3. 

22.  When  CaCl2-2H2O  is  heated,  water  is  liberated.  Write 
the  equation. 


23.  Balance  the  following  equations.  Check  each  one.  Do 
not  write  in  this  book. 

a)  Cu  +  S->Cu2S 

b)  P  +  C12->PC13 

c)  C  +  C02->CO 

d)  H2S  +  Pb(NO3)2  ->  HN03  +  PbS 

e)  CaO  +  G  ->  GaC2  +  CO 

f)  SiO2  +  G  -»  SiC  +  CO 

24.  Complete  and  balance  the  following  equations.  Check 
each  one.  Do  not  write  in  this  book. 

a)  H2O2  — >  e)  BaO2  +  H2SO4  — > 

b)  HgO->  f)  AgNO3  +  KCl-» 

c)  Fe  +  HCl->  g)  Mg  +  H2S04-» 

d)  CuO  +  H2-»  h)  Ca(OH)2  +  CO2  -> 

25.  Write  balanced  equations  for  the  following.  Name  each 
element  and  compound  that  appears  in  the  equation.  Check 
each  equation. 

a)  The  electrolysis  of  H2O. 

b)  The  laboratory  preparation  of  O2. 

c)  The  equation  representing  Priestley's  discovery  of  OL>. 

d)  The  laboratory  preparation  of  H2. 

e)  The  heating  of  Cu  in  air. 

f)  The  preparation  of  H2O2. 

g)  The  reduction  of  CuO^  by  H2. 

h)   The  heating  of  crystallized  washing  soda, 
i)   The  chemical  reaction  of  the  oxy hydrogen  torch, 
j)  The  action  of  Na  on  H2O. 
k)  The  passing  of  N2  over  hot  Mg. 
1)   The  reaction  between  H2O  and  CO2. 

26.  Is  there  anything  in  a  balanced  equation  which  did  not 
have  to  be  determined  by  experimentation? 



.  .  .  As  the  usefulness  and  accuracy 
of  chemistry  depend  entirely  upon 
the  determination  of  the  weights  of 
the  ingredients  and  products,  too 
much  precision  cannot  be  employed 
in  this  part  of  the  subject,  and  for 
this  purpose  we  must  be  provided 
with  good  instruments. 

Lavoisier,  1743-1794 

Measurement  in  chemistry.  Chemists  are  chiefly  concerned  in 
analyzing  substances  or  in  making  new  substances.  One  question 
they  always  ask  is  how  much  of  each  element  is  present  in  a  sub- 
stance. Their  measurements  of  quantity  must  be  very  exact.  Some- 
times their  answer  is  in  terms  of  the  volume  occupied  by  an  element. 
More  often  the  answer  they  want  is  in  terms  of  weight. 

Here  in  America,  in  everyday  affairs,  we  measure  the  weight  of 
substances  in  ounces  and  pounds.  Chemists  use  a  different  kind  of 
measurement.  They  weigh  substances  in  grams,  kilograms,  and  micro- 
grams.  They  also  measure  the  weight  of  atoms  in  terms  of  atomic 

All  measures  of  weight  are  comparisons.  When  you  say,  "My 
friend,  Charlie  Ross,  weighs  132  pounds,"  you  are  making  a  definite 
statement  that  anyone  can  understand.  But  what  is  a  pound?  You 
know  what  a  pound  of  butter  looks  like,  but  a  pound  by  itself 
doesn't  look  like  anything. 

Down  in  the  Bureau  of  Standards  in  Washington  is  a  carefully 
protected  metal  cylinder  called  the  prototype,  or  standard,  kilogram. 
Every  measurement  of  weight  made  in  our  nation  is  a  comparison 
with  that  particular  kilogram.  For  example,  the  pound  which  we 



use  in  everyday  measurements  of  weight  may  be  defined  as  0.4536 
of  the  standard  kilogram. 

Just  as  the  pound  is  divided  into  ounces,  the  kilogram  may  be 
divided  into  smaller  units  known  as  grams.  One  thousand  grams 
equal  one  kilogram.  The  gram  is  a  customary  unit  of  weight  in  chem- 
istry. When  chemists  speak  of  ten  grams  of  iron,  they  are  actually 
referring  to  an  amount  weighing  T^  of  the  standard  kilogram  in 

Chemists  also  use  a  Table  of  Atomic  Weights  in  which  one  atom  of 
each  element  has  a  particular  weight.  The  weight  assigned  to  each 
element,  like  all  measures  of  weight,  is  a  comparison.  How  was  this 
Table  of  Atomic  Weights  made? 

Dalton  determines  atomic  weights.  As  you  know,  Dal  ton  believed 
that  the  atoms  of  different  elements  have  different  weights.  Dalton 
knew  that  if  he  could  find  the  weights  of  the  atoms,  the  progress  of 
chemistry  would  be  speeded.  He  rcali/.ed  that  he  could  not  actually 
weigh  an  atom  of  an  element.  In  fact,  it  took  more  than  100  years 
from  the  time  of  Dalton's  experiments  until  accurate  methods  and 
precise  instruments  made  it  possible  to  determine  the  actual  weights 
of  atoms. 

However,  Dalton  knew  that  elements  combine  according  to  fixed 
ratios  by  weight.  For  example,  22.997  grams  of  sodium  combine  with 
79.916  grams  of  bromine  to  form  sodium  bromide  —  a  ratio  of  ap- 
proximately 1  :  3.5.  The  two  elements  never  vary  in  the  ratio  of  their 
weights  in  forming  this  compound.  In  a  similar  manner,  all  elements 
combine  according  to  certain  specific  ratios  by  weight.  Dalton  be- 
lieved that  the  ratios  of  these  combining  weights  depended  upon  the 
weights  of  the  atoms  of  each  element.  He  believed  further  that,  by 
studying  the  ratios  of  the  weights  in  which  elements  combine,  he 
would  be  able  to  determine  the  relative  weights  of  the  atoms  of  the 

The  prototype  kilogram  No.  20 
is  a  platinum-iridium  cylinder 
39  mm.  in  diameter  and  39  mm. 
high.  It  is  kept  at  Washington, 
D.C.,  in  the  laboratories  of  the 
Bureau  of  Standards. 


He  selected  the  atom  of  hydrogen  for  his  standard  and  assigned  to 
it  the  atomic  weight  1.  By  choosing  the  lightest  element  for  a  stand- 
ard, he  made  sure  that  the  atomic  weight  of  each  of  the  other  ele- 
ments would  be  greater  than  one.  Then  in  a  series  of  experiments, 
he  found  how  the  weight  of  the  atoms  of  13  other  elements  com- 
pared with  his  standard.  He  found  that  an  atom  of  oxygen  weighed 
seven  times  as  much  as  an  atom  of  hydrogen  and  assigned  to  oxygen 
the  atomic  weight  of  7;  he  found  that  an  atom  of  phosphorus 
weighed  nine  times  as  much  as  an  atom  of  hydrogen  and  assigned  to 
phosphorus  the  atomic  weight  of  9.  These  weights  have  since  been 
found  in  error,  but  that  in  no  way  detracts  from  the  value  of  Dalton's 

Dalton  prepares  an  historic  table.  From  his  experimental  data, 
Dal  ton  prepared  a  list  of  the  14  elements  arranged  according  to  the 
increasing  relative  weight  of  their  atoms.  This  was  the  first  table  of 
atomic  weights.  It  was  first  made  public  on  October  21,  1803  "before 
a  select  group  of  nine  members  and  friends  in  the  rooms  of  the 
Literary  and  Philosophical  Society  of  Manchester."  Although  inac- 
curate, the  table  compiled  by  this  Quaker  schoolteacher  remains  a 
monument  to  his  foresight  and  intellectual  accomplishment.  His 
achievement  was  a  crucial  advance  in  chemistry  and  formed  the 
keystone  of  his  theory. 

What  is  the  atomic  weight  of  an  element?  Later  investigators  fol- 
lowed Dalton's  method,  but  used  oxygen  as  a  standard  rather  than 
hydrogen.  They  found  oxygen  a  better  choice  because  it  combines 
with  far  more  elements  than  hydrogen.  The  atomic  weight  of  oxygen 
was  given  the  whole  number  16.  Many  of  the  other  elements  have 
atomic  weights  that  are  whole  numbers.  The  weight  of  hydrogen 
remains  approximately  one  —  actually  1.0078  in  comparison  to 

The  modern  Table  of  Atomic  Weights  appears  on  the  opposite 
page.  For  practical  purposes,  since  the  atomic  weight  of  hydrogen  is 
approximately  one,  hydrogen  can  still  be  used  as  the  basis  of  com- 
parison. The  number  200.6  after  mercury  means  that  one  atom  of 
mercury  weighs  a  little  less  than  200  times  as  much  as  an  atom  of 
hydrogen.  It  weighs  a  little  more  than  12|  times  as  much  as  an  atom 
of  oxygen. 

The  atomic  weight  of  an  element  is  a  number  that  shows  the 
comparison  of  the  weight  of  one  of  its  atoms  to  the  weight  of  one 
atom  of  oxygen,  which  is  considered  to  be  16. 

Always  remember  that  atomic  weights  are  not  measured  weights 
like  an  ounce  or  a  gram.  They  are  merely  relative  weights.  The 




AT.  WT. 



AT.  WT. 



















































































































actual  measured  weight  of  each  of  the  atoms  is  an  extremely  small 
quantity,  inconvenient  to  use  in  most  calculations.  For  example, 
the  actual  weight  of  an  atom  of  oxygen  is  about  0.000,000,000,000,- 
000,000,000,026  gram. 

The  Table  of  Atomic  Weights  is  the  foundation  of  chemical 
mathematics.  Realizing  the  importance  of  accurate  atomic  weights, 
many  chemists  performed  a  tremendous  number  of  experiments  to 
make  the  Table  of  Atomic  Weights  as  free  from  error  as  possible. 
Chemists  of  all  countries  cooperated  in  this  huge  undertaking.  In 
our  own  country,  Theodore  W.  Richards  and  his  students  at  Harvard 
spent  almost  half  a  century  in  this  epoch-making  work.  National 
boundaries  were  forgotten  and  men  from  all  over  the  world  pooled 
the  results  of  their  experiments  to  give  us  our  present  Table  of 
Atomic  Weights.  Research  chemists,  industrial  chemists,  and  student 
chemists  all  depend  upon  the  Table  of  Atomic  Weights  in  making 
their  chemical  calculations. 

Solving  type  problems  in  chemistry.  Most  of  the  problems  met  in 
elementary  chemistry  can  be  grouped  conveniently  into  five  types. 
The  first  two,  and  one  variety  of  the  third  are  described  below.  Two 
additipnal  varieties  of  the  third  type  are  discussed  in  Chapter  19 
and  two  more  types,  more  complex  and  less  frequently  met,  are 
discussed  on  pages  636—641. 



With  a  thorough  understanding  of  the  type  problems  discussed 
in  this  book,  you  should  have  no  trouble  in  solving  practically  all 
the  common  chemical  problems.  Frequent  reflective  practice  is,  of 
course,  necessary  for  mastery.  Hence  problems  of  various  types  are 
included  in  the  questions  at  the  end  of  each  of  the  remaining 

The  meaning  of  symbols.  A  knowledge  of  the  meaning  of  chemical 
symbols  is  essential  in  solving  the  various  types  of  chemical  problems. 
The  symbol  for  an  element,  like  K,  does  three  different  jobs.  First, 
it  may  be  used  to  name  the  element.  Second,  it  may  be  used  to  mean 
one  atom  of  that  element.  Third,  it  may  stand  for  one  atomic  weight 
of  the  element.  Thus  K  is  the  symbol  for  potassium,  for  one  atom  of 
potassium,  and  for  39  units  of  potassium  in  any  system  of  measuring 
weight  (as  39  grams,  39  ounces,  39  pounds) . 

The  chemical  symbols  for  a  formula  like  KCl  also  do  three  jobs. 
First,  they  represent  the  name  of  the  compound,  potassium  chloride. 
Second,  they  represent  one  molecule  of  potassium  chloride.  Third, 
they  represent  one  molecular  weight  of  the  compound.  All  of  the 
type  problems  involve  the  use  of  both  atomic  iveight  and  molecular 



Molecular  weight  is  the  ratio  of  the  weight  of  one  molecule 
of  a  compound  to  the  atomic  weight  of  oxygen  (16) .  Like 
atomic  weight,  it  is  only  a  relative  weight.  The  molecular 
weight  of  a  compound  is  obtained  by  adding  together  the 
atomic  weights  of  each  of  the  atoms  in  one  molecule  of  the 


1.  Find  the  atomic  weights  of  the  elements  in  the  chart  on 
page  127.  Place  these  numbers  under  the  symbols  and  add 

*  Some  compounds  do  not  exist  as  molecules  and  therefore  cannot  have  a 
molecular  weight.  However,  the  formulas  of  these  compounds  are  written  in  the 
conventional  manner.  The  steps  given  in  this  discussion  for  finding  molecular 
weight  may  also  be  used  for  finding  the  formula  weight  of  the  non-molecular 
compounds.  For  the  purposes  of  this  book,  formula  weight  may  be  considered  the 
same  as  molecular  weight. 


EXAMPLE:  Find  the  molecular  weight  of  potassium  chloride,. 

K     a 

39  +  35.5  =  74.5,  the  mol.  wt.  of  KC1 

2.  When  an  element  is  followed  by  a  subscript,  be  sure  to 
multiply  the  atomic  weight  of  the  element  by  the  subscript. 

EXAMPLE:  Find  the  molecular  weight  of  magnesium  sulfate, 

Mg      S  04 

24  +  32  +  (16  X  4)  =  24  +  32  +  64 

=  120,  the  mol.  wt.  of  MgSO4 

3.  When  a  formula  is  preceded  by  a  numerical  coefficient, 
multiply  the  total  molecular  weight  by  the  coefficient  to  find 
the  relative  weight  (rel.  wt.) . 

EXAMPLE:  Find  the  relative  weight  of  two  molecules  of  mag- 
nesium carbonate,  2MgCO3. 

2(Mg     C          08) 

2  [24  +  12  +  (16  X  3)] 

2  (24  +  12  +  48) 

2  (84)  =  168,  the  rel.  wt.  of  2  molecules  of  MgCO3 

4.  When  a  radical  is  enclosed  in  parentheses  followed  by  a 
subscript,  multiply  the  sum  of  the  atomic  weights  of  all  the 
elements  of  the  radical  by  the  subscript. 

EXAMPLE:  Find  the  molecular  weight  of  calcium  bicarbonate, 
Ca  (HC03)2. 

Ca      (H      C  O8)2 

40  +  [1  +  12  +  (16  X  3)]2 

40  +  (1  +  12  +  48)2 

40 +(6 1)2 

40  +  122  =  162,  the  mol.  wt.  of  Ca  (HCO8)2 

5.  Water  of  crystallization  (see  page  68)   is  chemically  part 
of  certain  compounds  and  is  separated  from  the  rest  of  the 
formula  by  a  dot.  The  dot  stands  for  plus  and  is  not  to  be 
considered  a  multiplication  sign. 


EXAMPLE:  Find  the  molecular  weight  of  crystallized  copper 
sulfate,  CuSO4  •  5H2O. 

Cu      S  04      •     5H2         O 

64  +  32  +  (16  X  4)  +  5(1  X  2  +  16) 
64  +  32  +  64  +5(18) 

64  +  32  +  64  +  90  =  250,  the  mol.  wt.  of  crystallized  copper 


The  gram-molecular  weight  or  mole.  A  very  convenient  unit  in 
many  calculations  is  the  gram-molecular  weight  or  mole,  which  is 
used  in  chemical  equations  (see  pages  282-283)  and  in  preparing 
standard  solutions  (see  pages  207-208) .  A  mole  is  the  molecular 
weight  of  a  substance  expressed  in  grams.  For  example,  a  mole  of 
potassium  chloride  (see  Procedure  1)  is  74.5  grams. 


1.  Calculate  the  molecular  weight  of  (a)  KBr  and  (b)  Nal. 

2.  Find  the  molecular  weights  of  (a)  LiCl  and  (b)  ZnO. 

3.  What  is  the  weight  of  a  mole  of  K3PO4? 

4.  Find  the  molecular  weights  of   (a)  H2SO4,    (b)  CaCO3> 
and  (c)  BaSO3. 

5.  Calculate   the  molecular  weights  of    (a)  Cu  (HCO3)  0, 
(b)Ba(N03)2,  and   (c)  A12  (SO4)  ,. 

6.  Calculate  the  molecular  weight  of  Na2S2O3  •  5H2O  (com- 
monly called  hypo) . 

7.  What  is  the  weight  of  a  mole  of  gypsum,  CaSO4  •  2H2O? 

8.  Find  the  molecular  weight  of  plaster  of  Paris, 


9.  Thfe  formula  of  glauber  salt  is  Na2SO4  •  10H.,O.  Find 
the  relative  weight  of  three  molecules  of  this  substance. 

10.  Calculate  the  relative  weight  of  4Na2B4O7  •  10H2O. 


The  percentage  composition  of  a  compound  is  found  by 
computing  the  percentage  by  weight  of  each  different  element 
in  the  compound.  This  is  a  simple  percentage  problem. 

Procedure.  Divide  the  atomic  weight  of  each  element  by  the  molecu- 
lar weight  of  the  compound  and  multiply  the  fraction  thus 
obtained  by  100. 


EXAMPLE  A:  Find  the  percentage  composition  of  nitric  acid, 

H      N          O3 

1-1-  14  +  (16X3) 

1  +  14  +  48  =  63,  the  mol.  wt.  of  HNO3 
%  of  hydrogen 

at.wt.ofH  1  X  100 

mol.wtofHNO,  ~         63 

%  of  nitrogen 

x  100  -  22.2% 

mol.  wt.  of  HNO3      ---  63 

%  of  oxygen 

_      rel.wt.of30  48X100 


mol.  wt.  of  HN03  63 

Total  =  100.0% 

EXAMPLE  B:  Find  the  percentage  of  water  of  crystallization  in 
BaCl2  -  2H2O. 

Ba  C12  2H2          O 

137  +  (35.5  X  2)  +  2(1  X  2  +  16) 

137  +  71  +2(18) 

137  +  71  +  36  =  244,  the  mol.  wt.  of  BaCl2  2H2O. 

Percentage  =    rcl.  wt.  of  2H2O    x  1Q() 
of  water         mol.  wt.  of  crystal 

EXAMPLE  C:  Find  the  percentage  composition  of 
Na2SO4    10H2O. 

Na2  S  O4        •       10H2         O 

(23  X  2)  +  32  +  (16  X  4)  +  10(1  X  2  +  16) 
46  +  32  +  64  +  180  =  322,  the  mol.  wt.  of  the  compound. 

Now  find  the  total  atomic  weight  of  each  element  in  the 
compound,  thus: 

Sodium       -    2  atoms  -  23  X    2  =    46,  the  total  at.  wt.  of  Na 
Sulfur         =     1  atom  =    32,  the  total  at.  wt.  of  S 

Oxygen      =  14  atoms  =  16  X  14  =  224,  the  total  at.  wt.  of  O 
Hydrogen  •  20  atoms  =  20  X     1  =    20,  the  total  at.  wt.  of  H 
Sum  of  the  at.  wt.  =  322,  the  mol.  wt.  of  the 



%  of  sodium  =  -3  X  100  =  14.3% 
%  of  sulfur  =  ^  X  100  =  10.0% 
%  of  oxygen  -  |||  X  100  -  69.5% 

%  of  hydrogen  =  ~  X  100  =       6.2% 

Total  =  100.0% 

EXAMPLE  D:  Find  the  weight  of  iron  in  80  Ib.  of  an  ore  con- 
taining 90  percent  ferric  oxide,  Fe2O3. 

The  weight  of  ferric  oxide  in  the  ore  is  90  percent  of  80  Ib., 
or  72  Ib. 

Fe2  O3 

(56  X  2)  +  (16  X  3)  =  160  =  mol.  wt.  of  Fe2O3 
Percentage  of  iron  in  Fe2O3  =  H$  X  100  =  70% 
Therefore  70%  of  72  Ib.,  or  50.4  Ib.,  is  the  weight  of  the  iron 
in  the  ore. 


1.  Calculate    the    percentage    composition    of     (a)  water, 
(b)  hydrogen  peroxide,  and  (c)  mercuric  oxide. 

2.  Calculate    the    percentage    composition    of    (a)  H2CO3, 
(b)  N2O4,  and   (c)  CaSO4  -  2H2O. 

3.  Find  the  percentage  composition  of  chrome  alum,  * 

KCr(S04)2-  12H20. 

4.  Calculate  the  percentage  composition  of  crystallized  potas- 
sium ferrocyanide,  whose  formula  is  K4Fe  (CN)  6  •  3H2O. 

5.  Find  the  percentage  of  oxygen  in  a  compound  whose 
formula  is  NiSO4  -  (NH4)  2SO4  -  6H2O. 

6.  How  much  aluminum  can  be  obtained  from  100  Ib.  of  its 
cryolite  ore  which,  upon  analysis,  shows  the  presence  of  80 
percent  Na3AlF0? 


Because  the  symbol  of  an  element  and  the  formula  of  a  com- 
pound may  represent  definite  weights,  an  equation  also  may 
be  considered  to  represent  definite  weights  of  the  substances 
taking  part  in  the  reaction.  Thus,  2Ag  -f-  S  — »  Ag2S  may  be 
read,  216  grams  of  silver,  plus  32  grams  of  sulfur  yield  248 
grams  of  silver  sulfide.  Note  that  the  actual  weights  are  based 
on  the  atomic  weights. 

Problems  based  on  chemical  equations  may  be  broadly  di- 
vided into  three  groups:  A.  Straight- weight  problems; 
B.  Weight-volume  problems;  and  C.  Straight- volume  problems. 


Group  A  is  discussed  below.  Groups  B  and  C  are  discussed 
later  in  the  book. 


Straight-weight  problems  involve  finding  one  weight  in  an 
equation  when  another  is  given. 

EXAMPLE:  How  many  grams  of  calcium  carbonate  will  be 
formed  by  the  complete  reaction  between  222  g.  of  calcium 
hydroxide  and  carbon  dioxide? 


1.  Write  the  balanced  equation. 

Ca(OH)2  +  C02  -»  H20  +  CaCO3 

2.  Write  the  given  weight  over  its  formula.  Write  x  over  the 
formula  of  the  substance  whose  weight  is  to  be  found.  Cross 
out  all  other  formulas  in  the  equation. 

222  g.  x  g. 

Ca(OH)2  +^Q*,^H2QL+  CaCO8 

3.  Since  the  same  relationship  exists  between  the  actual 
weights  as  exists  between  the  molecular  weights  represented 
in  the  equation,  write  the  molecular  weights  of  the  substances 
involved  under  their  respective  formulas.  Do  not  ignore  any 

222  g.  x  g. 

Ca       (O      H)2  -»       Ca      C  O3 

40  +  (16  +  1)2  40  +  12  +  (16  X  3) 

40  +  (17X2)  40+  12  +  48 

74  100 

4.  Write    the   mathematical   equation    represented   by    the 
known  and  unknown  weights.  Solve  for  x. 

222  =  _£_ 
74  100 
74*  =  22,200 

x  =  300,  the  number  of  grams  of  CaCOs  produced 

Alternate  method.  We  can  avoid  the  use  of  an  equation  in- 
volving x  by  solving  the  problem  as  follows: 


wt.  of  substance  used 


mol.  wt.  of  sub- 
X  P  -      =  Answer 

mol.  wt.  of  substance  used       stance  formed 
™  X  100  =  300,  the  number  of  grams  of  CaCO3  produced. 


1.  How  much  magnesium   is  required   to  react  with  suffi- 
cient hydrochloric  acid  to  produce  10  g.  of  hydrogen? 

Mg  +  2HC1  -*  MgCl2  +  H2 

2.  434   g.    of   mercuric   oxide,    HgO,   were   decomposed   by 
heat.  How  much  mercury  was  formed? 

3.  How    much    potassium    chlorate    would    be    needed    to 
prepare  384  g.  of  oxygen? 

4.  How  much  hydrogen  would  be  needed  to  reduce  com- 
pletely  100  g.  of  cupric  oxide,  CuO? 

5.  11.5   g.   of  sodium   react   completely  with  water.   What 
weight  of  sodium  hydroxide  is  formed? 

6.  By    the   electrolysis   of   water    12    Ib.    of   hydrogen   were 
liberated.   What  weight   of   oxygen   was   formed   at   the   same 

7.  (a)  What  weight  of  magnesium  will  be  needed  to  react 
with    sulfuric    acid    to    produce    30    g.    of    MgSO4?     (b)  What 
weight  of  hydrogen  will  be  evolved? 

8.  After  heating,    10  g.   of  crystalli/ed  copper  sulfate  gave 
6.4  g.  of  the  anhydrous  salt,  CuSO4.  Calculate  the  number  of 
molecules  of  water  of  crystallization  in  the  original  compound. 

Let  x  represent  the  number  of  molecules  of  water  of  crystal- 

10  g. 

6.4  g. 

3.6  g.  (that  is,  10  -  6.4) 

Now  complete  the  problem. 

Standard   Oil   Company    (A'./.) 

Delicate  instruments  such 
as  these  are  used  in  micro- 
chemistry,  the  branch  of 
chemistry  which  involves 
handling  extremely  small 
quantities  of  matter. 


9.  After  being  heated,  10  g.  of  crystallized  washing  soda 
gave  3.71  g.  of  Na2CO3.  Calculate  the  number  of  molecules 
of  water  of  crystallization. 

10.  If  4  g.  of  crystallized  barium  chloride  lost  0.59  g.  upon 
being  heated  to  constant  weight,  find  the  formula  of  the 
crystalline  salt. 


Jaffe,  Bernard.  Chemical  Calculations.  World  Book  Co., 
Yonkers,  1947.  A  systematic  presentation  of  the  solution  of 
type  problems,  with  1000  problems  arranged  progressively 
according  to  lesson  assignments. 

Kendall,  James.  At  Home  among  the  Atoms.  D.  Appleton- 
Century  Co.,  New  York,  1932.  "A  Few  Figures"  explains  atomic 
weights  in  a  novel  way.  Tells  how  atomic  weights  are  found. 


1.  The  atomic  weight  of  an  element  is  a  number  that  rep- 
resents the  ratio  of  the  weight  of  1  of  its  atoms  to  the  weight 
of  1  atom  of  oxygen.  Atomic  weights  are  all  relative  weights. 

2.  The  molecular  weight  of  a  compound  is  the  ratio  ol  the 
weight  of  1  molecule  of  the  compound  to  the  atomic  weight  of 
oxygen  (16) . 

3.  A  chemical  symbol,  in  addition  to  representing  an  ele- 
ment, represents  one  atomic  weight  of  that  element. 

4.  A  mole  of  a  substance  is  its  molecular  weight  expressed 
in  grams. 


Group  A 

1.  (a)  Who  issued  the  first  table  of  atomic  weights?  (b) 
Why  was  it  later  decided  to  use  the  weight  of  the  oxygen  atom 
as  a  standard  instead  of  the  weight  of  the  hydrogen  atom? 

2.  Exactly  what  is  meant  by  saying  that  oxygen  has  an  at. 
wt.  of  16? 

3.  Find  the  mol.  wt.  of  the  compounds  that  have  the  fol- 
lowing   formulas:     (a)  cupric    acetate,    Cu  (C2H3O2)  2  •  H2O; 
(b)  chloroplatinic      acid,      H2PtCl2;       (c)  microcosmic      salt, 
HNaNH4PO4-4H2O.  Refer  to  Table  4  on  page  127. 


I    ... 

4.  Find  the  percentage  composition  of  each  of  the  follow- 
ing  compounds:     (a)  BaCO3,     (b)  KMnO4,    (c)  K4Fe  (CN)  6. 
Check  each  result. 

5.  Determine  the  percentage  composition  of  each  of  the 
following  compounds:   (a)  BaSO4;  (b)  KCN;   (c)    (NH4)  2CO,. 

6.  Calculate  the  percentage  of  H0  in  alum, 


7.  Find  the  percentage  of  water  of  crystallization  in 

Sr  (N03)  2  •  5H20. 

8.  A  ton  of  limestone,  CaCOa,  was  heated  in  a  lime  kiln 
until  all  of  it  was  changed  to  quicklime,  CaO.  The  equation 
for   this  reaction   is:    CaCO3   -»   CaO   -+-   CO2j.   How   much 
quicklime  was  formed?  To  answer  this  question,  first  decide 
what  type  ot  problem  this  is:  What  four  steps  have  you  learned 
to  take  in  solving  such  a  problem?  Solve  the  problem. 

9.  Find  the  weights  of    (a)  H2  and    (b)  ZnSO4  formed  by 
the  complete  reaction  of  130  g.  of  Zn  and  sufficient  H.,SOt. 

Group  B 

10.  320  g.  of  Fe^Oy,  on  being  reduced,  form  224  g.  of  Fe. 
What  is  the  at.  wt.  of  oxygen? 

11.  11.95  g.  of  lead  sulfide,  PbS,  will  produce  10.35  g.  of 
lead.  From  this  fact,  and  knowing  that  the  at.  wt.  of  S  is  32, 
calculate  the  at.  wt.  of  Pb. 

12.  Of  what  use  to  the  manufacturing  chemist  is  knowledge 
of  the  percentage  composition  of  a  compound?  Select  from 
the  chapter  on  Cu  an  ore  of  that  metal  and  illustrate. 

13.  Suppose  a  chemist  were  going  to  manufacture  HC1  from 
NaCl  and  H2SO4.  What  helpful  information  could  he  gain 
from  the  following  equation  representing  the  reaction  that 
would  occur?  2NaCl  +  H2SO4 >Na2SO4  +  2HC1 


1.  Consult  a  manufacturing  chemist  or  an  analytical  chemist 
and  discuss  with  him  how  he  uses  chemical  mathematics  in 
his  business  or  profession.  Report  your  findings  to  your  class. 

2.  Construct  a  large  graph  to  represent  the  percentage  of 
each  element  in  crystallized  washing  soda,  Na2CO3  •  10H2O. 
Use   different  colors   for  each   element.   Show   also   the   per- 
centage of  H«Q*8JP;this  compound. 



.  .  .  Search  for  the  truth  is  in  one 
way  hard  and  another  easy,  for  it  is 
evident   tliat   no   one   can   master  it 
;  fully  nor  miss  it  wholly.  But  each 

adds  a  little  to  our  knowledge  of  Na- 
ture and  from  all  the  facts  assembled 
there  arises  a  certain  grandeur. 


A  Swedish  druggist  discovers  chlorine.  One  of  the  scientists  who 
lived  and  worked  at  about  the  same  time  as  Priestley,  Cavendish,  and 
Lavoisier  was  a  Swedish  pharmacist  who  is  well  known  to  the  world 
of  chemistry.  He  not  only  prepared  oxygen  earlier  than  Priestley, 
but  also,  in  1774,  discovered  chlorine.  Scheele  is  the  only  great  chem- 
ist whose  whole  lifework  was  accomplished  behind  the  counter  or 
in  the  prescription  laboratory  of  one  drugstore  or  another! 

When  he  had  discovered  the  greenish-yellow  chlorine  gas  that 
made  his  nose  and  throat  sting  and  almost  blinded  him,  he  wrote  to 
a  friend:  "Oh  how  happy  I  am;  1  seldom  think  of  eating,  or  drinking, 
or  where  I  live;  I  scarcely  pay  attention  to  my  pharmaceutical  busi- 
ness. But  to  watch  new  phenomena,  this  is  my  consuming  interest." 
At  43,  as  a  result  of  prolonged  exposure  to  unhealthful  conditions  in 
his  crude  laboratory,  Scheele  died  —  a  martyr  to  the  rapidly  develop- 
ing and  expanding  science  of  experimental  chemistry. 

How  chlorine  is  prepared  in  the  laboratory.  Hie  method  used  by 
Scheele  to  prepare  chlorine  is  still  the  common  laboratory  method 
used  today.  Two  compounds  are  used  —  manganese  dioxide  and 
hydrochloric  acid.  When  this  mixure  is  heated,  chlorine  is  liberated 
from  the  hydrochloric  acid. 




Fig.  21.  Laboratory  preparation  of  chlorine.  What 
property  of  chlorine  makes  it  wise  to  pass  the  excess 
gas  into  water? 

Although  chlorine  gas  is  fairly  soluble  in  water,  it  may  be  col- 
lected by  water  displacement.  Since  it  is  heavier  than  air,  it  is  us- 
ually collected  by  displacing  air  (see  Fig.  21)  . 

The  concentrated  hydrochloric  acid  used  in  this  reaction  supplies 
the  chlorine.  In  this  reaction  the  manganese  dioxide  acts  as  an  oxi- 
dizing agent,  combining  with  the  hydrogen  of  the  acid  and  liberating 
free  chlorine.  The  equation  for  this  reaction  is: 

[4HjCl  +  MnJO^i   -»  2H2O  +  MnCl2  +  C12 

This  equation  may  be  considered  to  represent  two  reactions.  The 
first  is  of  the  double-replacement  type: 

MnO2  +  4HC1  -»  MnCl4  +  2H2O 

and  the  second,  a  simple  decomposition:  *. 

MnCl4  -»  MnCl2  +  C12 

Other  oxidizing  agents  such  as  potassium  chlorate  or  lead  diox- 
ide, PbO2,  may  be  used  instead  of  manganese  dioxide. 

The  physical  properties  of  chlorine.  Chlorine  is  a  greenish-yellow 
gas,  two  and  one-half  times  as  heavy  as  air.  It  is  fairly  soluble  in 
water,  forming  yellowish  chlorine  water  (one  volume  of  water  dis- 
solves about  two  volumes  of  chlorine  gas  under  normal  laboratory 
conditions) .  It  has  a  penetrating  odor  and  attacks  the  membranes  of 
the  nose,  throat,  and  lungs.  Inhaling  ammonia  or  alcohol  vapor 
counteracts  this  irritating  action  to  some  extent.  Chlorine  was  the 
first  gas  to  be  liquefied.  It  liquefies  at  about  —  34°C.  at  normal 
pressure.  Faraday,  who  was  the  first  to  liquefy  it,  wrote  to  a  friend 
in  1823,  "I  hope  to  be  able  to  reduce  many  other  gases  to  the  liquid 
form."  He  did. 



Chemical  properties  of  chlorine.  Chlorine  is  a  typical  nonmetal. 
It  has  a  valence  of  one,  that  is,  it  combines  with  monovalent  hydro- 
gen, atom  for  atom.  Chlorine  is  very  active  chemically.  It  unites 
with  nearly  all  other  elements,  forming  compounds  called  chlorides, 
just  as  oxygen  forms  oxides.  Thus,  a  bit  of  sodium  reacts  brilliantly 
when  heated  with  moist  chlorine. 

Here  is  a  thrilling  example  of  the  marvels  of  chemistry.  Sodium,  a 
silvery,  poisonous  solid,  unites  with  chlorine,  a  greenish-yellow, 
poisonous  gas,  producing  a  white  solid,  common  salt,  which  is  essen- 
tial in  the  diet  of  both  man  and  all  animals. 

2Na  +  C12  ->  2NaCl 

Chlorine  has  a  strong  attraction  for  hydrogen.  When  hydrogen 
and  chlorine  are  mixed  and  exposed  to  strong  light  or  ignited  by  a 
spark,  they  combine  with  explosive  violence,  forming  hydrogen 
chloride  gas.  The  equation  for  this  reaction  is: 


If  a  jet  of  hydrogen  burning  in  air  is  thrust  into  a  jar  of  chlorine, 
it  will  continue  to  burn,  giving  off  hydrogen  chloride  as  the  product 
of  combustion. 

So  powerful  is  the  attraction  of  chlorine  for  hydrogen  that  it  will 
tear  hydrogen  away  from  some  of  its  compounds.  Thus,  when  tur- 
pentine, a  compound  of  carbon  and  hydrogen,  is  poured  over  a  piece 
of  filter  paper  which  is  then  thrust  into  a  bottle  of  chlorine,  a  flash 
of  light  occurs  and  a  black  powder  is  formed.  The  black  powder  is 
the  free  carbon  which  remains  after  the  hydrogen  from  the  turpen- 
tine has  combined  with  the  chlorine,  forming  hydrogen  chloride. 
Chlorine  will  combine  with  water,  liberating  oxygen. 

In  this  industrial  pla 
chlorine  is  being  man 
factored  in  mercury  cells. 


How  chlorine  is  prepared  for  industrial  use.  The  electrolysis  of 
brine,  a  water  solution  of  sodium  chloride,  is  the  source  of  most  of 
the  chlorine  used  today.  The  electric  current  liberates  free  chlorine 
at  the  carbon  anode.  At  the  cathode  the  sodium  liberated  reacts 
immediately  with  the  water,  forming  free  hydrogen  and  sodium 

2NaCl  ->  C12  +  2Na 
2Na  +  2H2O  -*  2NaOH  +  H2 

The  combined  equation,  then,  is: 

2NaCl  +  2H2O  -»  C12 1  +  H2  \  +  2NaOH 

The  chlorine  gas  is  dried  by  passing  it  through  concentrated  sul- 
luric  acid,  and  then  liquefied.  The  yellow,  liquid  chlorine  is  stored 
in  steel  cylinders,  each  holding  from  one  to  300  pounds  of  chlorine 
free  from  water  vapor.  This  process,  by  which  three  valuable  prod- 
ucts—  chlorine,  hydrogen,  and  sodium  hydroxide  —  are  formed 
from  two  low-cost,  plentiful  compounds  by  means  of  an  economical 
outlay  of  electric  energy,  is  described  in  greater  detail  on  page  212 
under  the  industrial  preparation  of  sodium  hydroxide.  Some  chlo- 
rine is  also  obtained  by  the  electrolysis  of  molten  NaCl  (see  page 

The  scientist  serves  humanity.  Soon  after  the  discovery  of  chlorine, 
Berthollet  hit  upon  the  idea  of  using  the  bleaching  action  of  chlorine 
(which  Scheele  had  noticed)  industrially.  He  declined  to  patent  his 
process  or  make  any  profit  from  it,  but  instead  turned  it  over  to  the 
French  government.  This  action  of  Berthollet  is  characteristic  of 
many  scientists,  who  believe  that  because  they  were  freely  helped 
toward  making  their  discoveries  by  the  work  of  those  who  preceded 
them,  so  they  should  freely  pass  on  the  results  of  their  own  labors. 

Fig.  22.  Chlorine  bleaches  indirectly.  In  the  presence  of  sunlight 
chlorine  reacts  with  water,  liberating  nascent  oxygen. 

chlorine  water  chlorine  wat^r  ir]^iWi^)t 

oxygen  gas- 


Patents  on  chemical  processes  are  sometimes  turned  over  by  their 
discoverers  to  foundations,  which  license  them  to  manufacturers. 
These  foundations  receive  fees  or  royalties  for  the  use  of  the  process. 
These  funds  may  be  used  by  the  foundations  in  making  further 
research  possible. 

Chlorine  bleaches  paper  and  textiles.  "Dephlogisticated  marine 
acid  air,"  as  chlorine  was  called  before  Davy  showed  it  to  be  an 
element,  was  a  chemical  curiosity  in  1774.  Today,  it  is  an  indus- 
trial necessity,  and  more  than  two  million  tons  of  it  are  produced  in 
this  country  annually. 

The  chief  use  of  chlorine  is  in  the  bleaching,  or  decolorizing,  of 
textiles,  chiefly  cotton  and  linen,  and  of  wood  pulp.  It  cannot  be 
used  for  bleaching  silk  or  wool  because  it  destroys  their  fibers.  Chlo- 
rine bleaches  indirectly,  by  oxidation.  Oxygen,  liberated  by  the  ac- 
tion of  chlorine  on  water,  combines  chemically  with  certain  coloring 
matters  and  dyes,  which,  upon  oxidation,  become  colorless.  This  re- 
action is: 

???  +!?•  °  ~* 2HC1  +  °  t 

Chlorine  is  used  as  a  bleaching  agent  either  in  the  free  condition, 
in  chlorine  water,  or  in  some  unstable  chlorine  compounds,  such  as 
bleaching  powder,  CaOCl2,  and  calcium  hypochlorite,  Ca  (CIO)  2. 
Laundry  bleach,  used  extensively  today,  is  made  by  adding  liquid 
chlorine  to  a  very  cold  solution  of  sodium  hydroxide.  This  product 
is  known  as  Javelle  water  and  contains  sodium  hypochlorite,  NaCIO, 
a  salt  of  hypochlorous  acid,  HC1O. 

2NaOH  +  C12  -*  NaCIO  +  NaCl  +  H2O 

Sodium  hypochlorite  decomposes  easily,  liberating  atomic  oxygen. 

Dilute  solutions  of  sodium  hypochlorite,  in  strengths  ranging  from 
four  to  six  percent,  are  marketed  as  "Clorox,"  "Rose-X,"  and  a  num- 
ber of  other  trademarks.  Such  household  bleaches  are  used  very 
widely  in  home  laundering  and  in  removing  stains.  In  bleaching  with 
chlorine,  care  must  be  taken  not  to  expose  the  materials  to  the  action 
of  the  chlorine  for  too  long  a  time,  since  continued  action  will  make 
the  fibers  very  weak.  After  the  bleaching  agent  has  been  used,  the 
fabrics  should  be  thoroughly  rinsed  in  fresh  water  to  remove  all 
traces  of  the  bleach.  In  industrial  bleaching,  after  removal  from  the 
bleaching  tank,  an  antichlor,  such  as  sodium  thiosulfate  (Na2S2O.<) , 
commonly  called  hypo,  is  added  to  remove  the  excess  chlorine,  and 
then  the  material  is  thoroughly  rinsed  in  running  water. 


Chlorine  saves  lives.  Although  it  consumes  only  about  ten  percent 
of  normal  production,  a  second  very  important  use  of  chlorine  is  in 
the  purification  of  water.  When  chlorine  is  added  to  water,  atomic 
oxygen,  liberated  from  the  water  by  chlorine,  reacts  with  the  bacteria 
present  and  kills  them  by  oxidation.  Only  three  pounds  of  chlorine 
are  used  for  each  million  gallons  of  water.  In  the  United  States  more 
than  75  percent  of  all  drinking  water  is  treated  with  chlorine.  As  a 
result,  the  death  rate  from  typhoid  fever,  a  disease  caused  by  typhoid 
bacilli  which  may  be  present  in  drinking  water,  has  been  cut  down 
considerably.  This  treatment  also  kills  algae  and  other  low  plant 
and  animal  life.  For  field  use,  in  areas  where  pure  water  is  not  avail- 
able, explorers,  scouts,  and  others  carry  tablets  of  a  chlorine-produc- 
ing substance  such  as  "halazone"  which  contains  Ca  (CIO)  2. 

Chlorine  is  also  used  as  an  antiseptic  and  disinfectant.  A  substance 
that  checks  the  growth  of  bacteria  is  called  an  antiseptic,  while  a 
germicide  kills  the  bacteria  outright.  A  substance  that  either  kills  or 
checks  the  growth  of  bacteria  is  called  a  disinfectant.  Since  much  ill- 
ness is  caused  by  poisons,  or  toxins,  formed  by  bacteria,  it  is  impera- 
tive that  these  bacteria  be  killed  or  at  least  prevented  from  mul- 

"Zonite,"  a  trade  preparation  containing  NaCIO,  is  used  as  a 
general  household  disinfectant.  Because  chlorine  disinfectants  if  used 
improperly  may  destroy  body  tissue,  such  preparations  should  be  used 
with  great  care.  Several  widely  used  insecticides  are  chlorine  com- 
pounds (see  page  540) . 

And  chlorine  destroys  human  life,  too.  Chlorine  played  a  double 
role  during  World  War  I.  While  chlorine  disinfectants  and  water 
chlorination  were  saving  thousands  of  lives,  free  chlorine  in  the 
form  of  clouds  of  poisonous  gas  was  choking  out  many  other 
lives.  Chlorine,  and  later  phosgene,  COCL,  and  mustard  gas, 
(C1CH2CH2)  2S,  caused  terrible  suffering  in  World  War  I,  even 
though  the  use  of  poisonous  gases  had  been  "outlawed"  by  a  confer- 
ence at  The  Hague  in  1907. 

How  to  test  for  the  presence  of  a  chloride.  Free  chlorine  may  be 
recognized  by  its  characteristic  color  and  odor,  but  compounds  of 
chlorine  cannot  be  identified  so  easily.  Because  chlorides  are  presejat 
in  so  many  common  substances,  a  simple  test  for  a  chloride  is  de- 

All  chlorides  are  soluble  in  water,  with  the  exception  of  the 
chlorides  of  silver,  mercury,  and  lead.  Because  silver  chloride  is  in- 
soluble in  water,  chlorides  are  recognized  by  their  reaction  with  a 
solution  of  silver  nitrate.  When  a  solution  of  silver  nitrate  is  added 


silver  nitrate 

Fig.  23.  Testing  an  unknown  solution  for 
the  presence  of  a  chloride.  What  will  hap- 
pen to  the  white  precipitate  on  exposure 
to  light  if  a  chloride  is  present? 

to  a  solution  of  a  chloride,  a  white  insoluble  substance,  a  precipitate, 
forms.  On  exposure  to  light,  the  color  of  this  precipitate  changes 
gradually  to  purple  and  then  to  black. 

AgNO3  +  NaCl  -»  NaNO3  +  AgCl  | 

This  color  change,  when  silver  nitrate  is  added  to  an  unknown 
solution,  suggests  that  probably  a  chloride  is  present.  But  some 
substances  that  are  not  chlorides  form  similar  compounds.  Therefore, 
the  formation  of  such  a  precipitate  on  the  addition  of  silver  nitrate 
to  a  solution  is  not  an  entirely  reliable  test  for  a  chloride. 

A  chemical  test  for  a  substance  has  one  important  requirement  — 
it  must  be  specific,  that  is,  no  other  substance  will  react  to  the  test  in 
the  same  way  as  the  substance  for  which  the  test  was  designed.  For- 
tunately, silver  chloride  is  insoluble  in  dilute  nitric  acid;  other  sub- 
stances that  might  at  first  be  mistaken  for  it  are  soluble  in  dilute 
nitric  acid  and  disappear  at  once  when  this  acid  is  added.  The  addi- 
tion of  nitric  acid,  therefore,  is  the  final  step  in  this  test,  distinguish- 
ing the  chloride  from  other  compounds. 

An  element  is  known  by  the  company  it  keeps.  The  halogens  (salt 
formers)  are  a  group  of  elements  that  resemble  one  another  chemi- 
cally, and  whose  physical  properties  differ  from  one  another  in 
regular  gradation,  as  shown  in  Table  5.  Such  a  group  of  elements 
is  called  a  chemical  family.  The  members  of  the  halogen  family  are 
fluorine,  chlorine,  bromine,  iodine,  and  astatine,  a  radioactive  ele- 
ment. Table  5  shows  the  relationship  among  the  members  of  the 
halogen  family. 

Making  bromine  in  the  laboratory.  Like  chlorine,  bromine  is  pre- 
pared by  the  oxidation  of  its  hydrogen  compound  by  manganese  diox- 
ide. A  mixture  of  potassium  bromide,  sulfuric  acid,  and  manganese 
dioxide  is  heated  in  a  test  tube  (see  Fig.  24) .  The  H2SO4  reacts 
with  the  KBr,  forming  HBr,  which  is  then  oxidized  by  the  MnO2. 
Although  at  room  temperature  bromine  is  a  liquid,  it  is  liberated  as 























Greenish  - 




BROMINE,  Br2  1 







IODINE,  I2       1 



Purplish  - 


Least  soluble 
of  halogens 


a  brownish  vapor  at  the  temperature  of  the  experiment.  As  this 
brownish  vapor  passes  into  the  water,  some  of  the  bromine  dissolves; 
the  rest  collects  as  a  layer  of  bromine  under  the  water.  This  method 
of  collecting  pure  bromine  by  distillation  is  relatively  safe. 

2KBr  +  2H2SO4  +  MnO2  ->  K2SO4  +  MnSO4  +  2H2O  +  Br2 

Great  care  should  be  taken  in  working  with  bromine,  because  it  is 
poisonous  and  attacks  the  skin,  causing  severe  burns.  Particular  care 
should  be  taken  to  protect  the  eyes  from  bromine  vapor. 

Taking  bromine  from  the  sea.  Most  bromine  is  extracted  from  the 
minute  percentage  (0.0065%)  of  bromides  present  in  sea  water. 
Free  chlorine  replaces  the  bromine  of  the  bromides.  Some  of  our 
bromine  is  also  obtained  from  the  bromides  found  in  salt  wells  and 
salt  lakes.  The  principal  chemical  reaction  is: 

MgBr2  +  C12  -» MgCl2  +  Br2 

Bromine  helps  engine  efficiency.  Much  of  the  bromine  produced 
in  the  United  States  is  used  in  the  manufacture  of  "Ethyl  fluid," 
an  anti-knock  mixture  composed  of  ethylene  dibromide  and  tetra- 
ethyl  lead  (TEL) .  Large  quantities  of  bromine  are  also  used  in 
making  silver  bromide,  the  light-sensitive  chemical  that  forms  the 
most  important  part  of  the  coating  of  photographic  films.  Bromine 
and  bromine  compounds  are  used  also  in  making  tear-gas  bombs. 
Bromine  is  used  in  appreciable  quantities  as  an  oxidizing  agent  in 
the  manufacture  of  certain  dyes  and  drugs. 

Potassium  bromide  or  sodium  bromide  acts  as  a  depressant  on  the 
central  nervous  system.  Their  action  is  followed  by  drowsiness  and 
even  sleep.  Such  chemicals  are  called  sedatives.  They  are  used  in  the 













Forms  HF 


H20  +  F2 
-*2HF  +  0 




Forms  HCI 


H20  +  CI2 
-*2HCI  +  0 




Forms  HBr 


-*  2HBr  +  O 

Less  than 



Forms  HI 


I2  +  H20 

-»  2HI  +  0 




treatment  of  insomnia  and  asthma  and  are  frequently  found  in  head- 
ache and  sleeping  powders.  Heavy,  continuous  doses  of  bromides  may 
have  harmful  effects  on  the  body.  Bromides  should  be  used  only  on 
the  advice  of  a  physician. 

Methyl  bromide,  CH;5Br,  is  used  widely  in  commercial  and  indus- 
trial fumigation  to  kill  insects  and  other  low  forms  of  plant  and  ani- 
mal life.  It  is  used  for  this  purpose  in  boxcars,  warehouses,  and  food 
processing  and  packaging  plants. 

How  iodine  is  prepared  in  the  laboratory.  Like  chlorine  and 
bromine,  iodine  is  prepared  by  the  oxidation  of  its  hydrogen  com- 
pound by  manganese  dioxide.  A  mixture  of  potassium  iodide,  sulfuric 
acid,  and  manganese  dioxide  is  heated  in  a  test  tube  as  shown  in 
Fig.  25. 

SKI  +  2H2S04  +  Mn02  -»  K2SO4  +  MnSO4  +  2H2O  -f  I2 

The  violet  vapor  of  iodine,  which  is  produced  at  the  temperature 
of  the  reaction,  condenses,  forming  purplish-black  crystals  on  the  in- 
side of  the  test  tube.  This  process  of  collecting  iodine  is  called  sub- 

Fig.  24.  Preparation  of  bromine  in 
the  laboratory.  Why  is  bromine  not 
collected  in  the  same  way  as  chlo- 





























-  187°C 


Forms  HF 


H20  +  F2 
-*  2HF  +  O 












Forms  HCI 


H20  +  CI2 
-*•  2HCI  +  O 




BROMINE,  Br2  1 








Forms  HBr 


Br2  +  H2O 
-*  2HBr  +  O 

Less  than 



IODINE,  I2        f 



Purplish  - 


Least  soluble 
of  halogens 

200°  C 


Forms  HI 


I2  +  H2O 
-*  2HI  +  0 




a  brownish  vapor  at  the  temperature  of  the  experiment.  As  this 
brownish  vapor  passes  into  the  water,  some  of  the  bromine  dissolves; 
the  rest  collects  as  a  layer  of  bromine  under  the  water.  This  method 
of  collecting  pure  bromine  by  distillation  is  relatively  safe. 

2KBr  +  2H2SO4  +  MnO2  -»  K2SO4  +  MnSO4  +  2H2O  +  Br2 

Great  care  should  be  taken  in  working  with  bromine,  because  it  is 
poisonous  and  attacks  the  skin,  causing  severe  burns.  Particular  care 
should  be  taken  to  protect  the  eyes  from  bromine  vapor. 

Taking  bromine  from  the  sea.  Most  bromine  is  extracted  from  the 
minute  percentage  (0.0065%)  of  bromides  present  in  sea  water. 
Free  chlorine  replaces  the  bromine  of  the  bromides.  Some  of  our 
bromine  is  also  obtained  from  the  bromides  found  in  salt  wells  and 
salt  lakes.  The  principal  chemical  reaction  is: 

MgBr2  +  C12  — >  MgCl2  +  Br2 

Bromine  helps  engine  efficiency.  Much  of  the  bromine  produced 
in  the  United  States  is  used  in  the  manufacture  of  "Ethyl  fluid," 
an  anti-knock  mixture  composed  of  ethylene  dibromide  and  tetra- 
ethyl  lead  (TEL) .  Large  quantities  of  bromine  are  also  used  in 
making  silver  bromide,  the  light-sensitive  chemical  that  forms  the 
most  important  part  of  the  coating  of  photographic  films.  Bromine 
and  bromine  compounds  are  used  also  in  making  tear-gas  bombs. 
Bromine  is  used  in  appreciable  quantities  as  an  oxidizing  agent  in 
the  manufacture  of  certain  dyes  and  drugs. 

Potassium  bromide  or  sodium  bromide  acts  as  a  depressant  on  the 
central  nervous  system.  Their  action  is  followed  by  drowsiness  and 
even  sleep.  Such  chemicals  are  called  sedatives.  They  are  used  in  the 



treatment  of  insomnia  and  asthma  and  are  frequently  found  in  head- 
ache and  sleeping  powders.  Heavy,  continuous  doses  of  bromides  may 
have  harmful  effects  on  the  body.  Bromides  should  be  used  only  on 
the  advice  of  a  physician. 

Methyl  bromide,  CH,,Br,  is  used  widely  in  commercial  and  indus- 
trial fumigation  to  kill  insects  and  other  low  forms  of  plant  and  ani- 
mal life.  It  is  used  for  this  purpose  in  boxcars,  warehouses,  and  food 
processing  and  packaging  plants. 

How  iodine  is  prepared  in  the  laboratory.  Like  chlorine  and 
bromine,  iodine  is  prepared  by  the  oxidation  of  its  hydrogen  com- 
pound by  manganese  dioxide.  A  mixture  of  potassium  iodide,  sulfuric 
acid,  and  manganese  dioxide  is  heated  in  a  test  tube  as  shown  in 
Fig.  25. 

2KI  +  2H2SO4  +  MnO2  ->  K2SO4  +  MnSO4  +  2H2O  +  I2 

The  violet  vapor  of  iodine,  which  is  produced  at  the  temperature 
of  the  reaction,  condenses,  forming  purplish-black  crystals  on  the  in- 
side of  the  test  tube.  This  process  of  collecting  iodine  is  called  sub- 

KBr  +  MnO2 

Fig.  24.  Preparation  of  bromine  in 
the  laboratory.  Why  is  bromine  not 
collected  in  the  same  way  as  chlo- 


iHk"  bromine 

Dow  Chemical  Company 

Methyl  bromide  used  as  a  fumigant  destroys  all  grain  insects. 

limation.  A  substance  is  said  to  sublime  when  it  passes  directly  from 
the  solid  state  to  the  gaseous  state  and  then  condenses  back  to  the 
solid  state  without  passing  through  the  liquid  state.  Camphor1  is  an- 
other substance  that  sublimes. 

The  industrial  preparation  of  iodine.  About  90%  of  our  iodine 
is  obtained  from  brine  that  comes  up  with  the  oil  in  California  oil 
fields.  This  brine  contains  Nal  and  MgL.  Chlorine  is  passed  through 
the  brine  and  replaces  the  iodine.  The  iodine  is  recovered  by  adsorp- 
tion on  activated  carbon.  (Adsorption  is  the  clinging  of  molecules  of 
one  substance  to  the  surface  of  another  —  see  page  327.)  The  princi- 
pal reaction  is: 

MgI2  +  C12  ->  MgCl2  +  I2 

The  rest  of  our  iodine  comes  from  NaIO;5,  found  as  an  impurity 
in  Chile  saltpeter,  NaNO,. 

Iodine,  too,  saves  lives.  The  chief  use  of  iodine  is  in  the  prepara- 
tion of  tincture  of  iodine,  a  two  percent  solution  of  iodine  and  po- 
tassium iodide  in  ethyl  alcohol,  which  is  an  excellent  antiseptic.  As 
silver  iodide,  iodine  is  used  to  some  extent  in  photography  together 
with  silver  bromide.  It  is  used  also  in  the  manufacture  of  iodoform, 
CHI:1,  a  yellow  powder  used  as  an  antiseptic,  and  in  the  manufacture 
of  "Aristol,"  an  improvement  over  iodoform.  Iodine  is  also  used  in 
the  production  of  certain  dyes  and  methyl  iodide,  CHJ. 




Is  iodine  necessary  to  health?  Iodine  is  an  important  constituent 
of  the  human  body.  There  is  a  definite  relation  between  the  presence 
of  iodine  in  the  thyroid  gland  and  the  prevalence  of  certain  disorders. 
The  thyroid  gland,  located  in  the  neck,  secretes  a  compound  called 
thyroxin,  containing  about  65  percent  iodine,  which  helps  to  regulate 
the  rate  of  oxidation  in  the  body. 

When  the  thyroid  gland  receives  too  little  iodine,  goiter,  an  en- 
largement of  the  thyroid,  results,  caused  apparently  by  the  attempt  of 
the  gland  to  increase  its  size  in  order  to  produce  more  thyroxin.  To 
offset  this  deficiency,  iodides  may  be  added  to  drinking  water  or 
about  0.02  percent  of  sodium  iodide  added  to  so-called  iodized  salt. 

Extreme  underactivity  ot  the  thyroid  gland  in  newborn  babies  and 
young  children  may  result  in  cretinism  (kre'tm-izm)  —  misshapen 
dwarfishness,  low  mentality,  sluggishness,  dullness,  slow  heart  action. 
Synthetic  thyroxin  is  used  in  the  treatment  of  this  thyroid  disorder. 
Overactivity  of  the  thyroid  gland  often  produces  the  opposite  effect  — 
the  thin,  nervous,  highly  energetic  person,  whose  movements  are 
quick,  and  whose  heart  action  is  rapid.  See  page  34  for  a  discussion 
of  basal  metabolism  tests  in  diagnosis  of  thyroid  disturbances. 

Iodine  is  also  necessary  to  other  forms  of  animal  life.  Large  quan- 
tities of  iodides  are  added  to  commercial  feeds  for  chickens,  cattle, 
dogs,  cats,  and  other  animals,  and  to  fertilizers  for  forage  crops. 

Replacement  power  of  the  halogens.  If  free  chlorine  is  added  to  a 
solution  of  a  bromide  or  an  iodide,  free  bromine  or  free  iodine  is 
liberated.  Free  chlorine  replaces  the  two  less  active  halogens. 

2KBr  +  C12  ->  2KC1  +  Br2 
2NaI  +  C12  ->  2NaCl  +  I2 

The  addition  of  free  bromine  to  a  solution  of  an  iodide  releases 
free  iodine. 

2KI  +  Br2  -*  2KBr  +  I2 

test  tube  containing          I 

'        H2SO4 


Fig.   25.   Laboratory   preparation   of 

iodine  crystals 


However,  the  addition  of  iodine  to  a  solution  of  either  a  bromide  or 
a  chloride  has  no  effect,  for  the  less  active  halogen  will  not  replace 
the  more  active  halogen  from  its  compound.  As  mentioned  earlier, 
the  commercial  preparation  of  bromine  depends  on  the  replacement 
power  of  chlorine. 

How  we  test  for  the  presence  of  bromides  and  iodides.  Many  uses 
are  made  of  the  replacement  power  of  the  halogens.  The  tests  for 
bromides  and  iodides  are  based  on  it.  Chlorine  water  is  added  to  a 
solution  of  the  unknown  salt,  and  a  few  drops  of  carbon  disulfide, 
CS2,  which  is  not  soluble  in  water,  are  also  added.  If,  after  shaking 
the  mixture,  the  carbon  disulfide  settles  out  as  a  distinct  layer  with 
a  brownish-red  coloration,  then  the  original  salt  was  a  bromide,  the 
free  bromine  liberated  coloring  the  carbon  disulfide.  If  the  carbon 
disulfide  acquires  a  purple  coloration,  the  original  salt  was  an  iodide. 

Carbon  disulfide  is  used  because  free  bromine  and  iodine  are  much 
more  soluble  in  this  liquid  than  in  water.  Hence,  most  of  the  liber- 
ated bromine  and  iodine  dissolve  in  the  carbon  disulfide,  thus  color- 
ing it  much  more  than  they  would  color  water.  Carbon  disulfide,  be- 
cause it  is  a  better  solvent,  will  extract  any  bromine  or  iodine  from 
the  water  solution.  This  process  of  separation,  frequently  used  in 
industry,  is  called  extraction  by  partition. 

Fluorine,  the  most  active  of  all  the  elements.  Fluorine  was  not 
isolated  until  1886.  Because  of  its  extreme  chemical  activity,  which 
causes  it  to  unite  violently  with  metals,  glass,  porcelain,  and  water, 
its  separation  as  a  pure  element  was  a  very  difficult  undertaking. 
Finally,  Henri  Moissan  succeeded  by  liquefying  pure  hydrogen 
fluoride,  adding  some  potassium  fluoride,  and  at  a  temperature  of 
— 23°C.  passing  an  electric  current  through  the  mixture.  Fluorine 
was  liberated  at  the  anode. 

The  anode  used  industrially  today  is  made  of  graphite,  which  is 
not  attacked  by  this  pale  yellow  gas.  Fluorine  is  employed  in  making 
uranium  hexafluoride  used  in  atomic  energy  plants  where  many  of 
the  lubricants  are  chemically  inert  fluorocarbon  compounds.  The 
plastic,  "Teflon,"  is  another  fluorine  compound.  The  new  rat  and 
ground  squirrel  poison,  1080,  is  a  fluorine  compound,  and  sodium 
fluoride  is  used  in  some  insecticides. 

Fluorine  prevents  tooth  decay.  The  amount  of  tooth  decay,  or 
dental  caries,  has  been  found  to  vary  directly  with  the  quantities  of 
fluorides  present  in  the- local  water  supply.  Too  much  fluoride  pro- 
duces very  hard  but  mottled  teeth,  a  condition  in  which  the  enamel 
becomes  discolored  badly.  When  too  little  fluoride  is  present,  there 
is  much  tooth  decay.  Water  fluoridation  (about  one  part  NaF  per 



million  parts  water)  is  now  widely  practiced  to  protect  children  up 
to  about  the  age  of  12  while  tooth  enamel  is  being  formed. 

Fluorine  refrigerants.  A  refrigerant  is  a  substance  used  to  absorb 
heat  by  changing  from  a  liquid  to  a  gas.  In  refrigeration,  the  material 
from  which  heat  is  absorbed  is  cooled.  Almost  as  long  as  mechanical 
refrigerators  have  existed,  their  manufacturers  have  searched  for  bet- 
ter refrigerants.  Something  was  wrong  with  nearly  all  the  original 
refrigerants.  They  were  either  toxic,  flammable,  corrosive,  or  pos- 
sessed disagreeable  odors.  And  then  a  family  of  compounds  was  de- 
veloped, and  introduced  in  1931  under  the  trademark  "Freon." 
These  compounds,  produced  by  the  halogenation  of  simple  com- 
pounds of  carbon  and  hydrogen,  are  far  superior  to  sulfur  dioxide, 
ammonia,  ethyl  chloride,  and  methyl  chloride  as  refrigerants.  All  are 
practically  odorless,  nontoxic,  nonflammable,  and  noncorrosive. 

The  first  of  the  "Freons"  to  be  produced  was  dichlorodifluoro- 
methane,  CC1,F,,  and  a  later  one  was  dichloroteArafluoroethane, 
C2C12F4.  There  are  several  others,  each  with  slightly  different  proper- 
ties that  make  it  particularly  well  adapted  to  a  special  use.  Because  of 
their  extreme  volatility  and  the  speed  with  which  they  penetrate 
every  nook  and  cranny  of  a  confined  space,  the  "Freons"  are  used  as 
propellants  in  dispersing  insecticides.  Aerosol  bombs  used  for  killing- 
household  pests  usually  contain  an  insecticide  and  a  liquefied  "Freon" 
gas  under  pressure.  When  the  pressure  is  released,  the  expanding  gas 
quickly  distributes  the  insecticide  throughout  a  room  or  closet. 

This  electronic  device,  called  a  "sniffer/1  is  used  to  detect 
breaks  in  telephone  cables  through  which  moisture  might 
enter.  The  cable  to  be  tested  is  filled  with  a  "Freon1'  gas. 
Then  the  sniffer  is  pulled  along  the  cable  from  the  ground. 
Escaping  gas  activates  a  "FreorT'-sensitive  detector,  ringing 
a  bell. 

Bdl  Telephone  laboratories 




Consumer  Reports  (Consumers  Union  Reports)  ,  April,  May, 
June,  1939.  Consumers  Union  of  U.S.,  New  York.  Excellent 
reports  on  various  antiseptics. 

Harrow,  Benjamin.  Eminent  Chemists  of  Our  Time  (2nd 
ed.) .  D.  Van  Nostrand  Co.,  New  York,  1927.  Read  the  life  of 
Moissan  and  his  isolation  of  fluorine. 

Sanders,  Gardiner,  and  Wood.  "Chlorine  and  Caustic  Manu- 
facture/' Industrial  and  Engineering  Chemistry,  September, 
1953,  pp.  1824-1835.  Includes  history,  production  figures, 
photos,  and  diagrams  showing  how  diaphragm  and  mercury 
cells  work. 


1.  The  halogens  are  a  group  of  elements  that  resemble  one 
another  chemically,  and  whose  physical  properties  differ  in 
regular  gradation.  The  halogen  group  is  one  of  several  such 
groups  of  elements. 

2.  A  substance  that  passes  directly  from  the  solid  to  the 
gaseous  state  and  directly  from  the  gaseous  to  the  solid  state 
is  said  to  sublime. 

3.  The  ability  of  one  element  to  replace  another  in  that  ele- 
ment's compounds  is  widely  used  in  chemical  reactions.  The 
halogens  may  be  listed  according  to  their  replacement  power. 

4.  The  separation  of  a  substance  from  a  solution  containing 
that  substance  by  means  of  a  better  solvent  is  known  as  extrac- 
tion by  partition. 

5.  Like  Berthollet,  many  scientists  believe  they  should  freely 
pass  on  to  others  the  results  of  their  own  labors,  and  so  help 
all  humanity. 


Group  A 

1.  Who  discovered  C12,  and  in  what  year  was  it  isolated? 

2.  Make  a  labeled  diagram  showing  the  laboratory  prepara- 
tion of  C12. 

3.  What  is  the  function  of  MnO2  in  the  laboratory  prep- 
aration of  C12? 

4.  What  other  substance  might  be  used  instead  of  MnO2  in 
the  preparation  of  C12? 


•<5.  (a)  .Give  two  reasons  for  collecting  C12  by  the  displace- 
ment of  air.  (b)  Why  should  any  excess  of  C12  be  passed  into 

6.  Write  a  balanced  equation  representing  the  laboratory 
preparation  of  Cla. 

t  .  .  . 

7.  State  four  physical  properties  of  C12. 

8.  Faraday  liquefied  C12  in  1823.  It  was  the  first  gas  to  be 
liquefied.  Suggest  a  reason. 

9.  If  a  brightly  burning  paraffin  taper  is  inserted  in  a  bottle 
of  C12,  a  heavy  black  smoke  is  given  off.  Explain. 

10.  By  what  process  does  C12  bleach?  Explain. 

11.  (a)  Why  is  it  necessary   to  rinse  materials   that  have 
been  bleached  with  C12?  (b)  Why  cannot  C12  be  used  to  bleach 
silk  and  wool? 

12.  Because  of  what  chemical  property  are  bleaching  pow- 
der, CaOCl2,  and  Clorox,  containing  NaOCl,  able  to  bleach? 

13.  Find  the  percentage  of  C12  in  bleaching  powder. 

14.  (a)  Distinguish  between  the  terms  antiseptic  and  germi- 
cide, (b)  What  term  embraces  both? 

15.  Make  a  list  of  all  the  uses  of  C12  that  you  know. 

16.  Chloride  of  lime,  or  bleaching  powder,  has  the  formula 
CaOCl2.  It  is  made  from  Ca  (OH)  2  and  C12.  Write  the  equa- 
tion for  its  preparation. 

17.  Give  a  brief  account  of  the  preparation  of  some  laundry 

18.  Chlorine  played  a  double  role  in  wartime.  Explain. 

19.  How  much  CI2  can  be  prepared  by  the  action  of  348  g. 
of  MnO2  on  sufficient  concentrated  HC1? 

.  .  t  .  .  . 

"  1 

20.  (a)  Describe  fully  the  test  for  a  chloride,  (b)  Write  the 
equation  for  the  reaction  that  takes  place. 

21.  Illustrate  the  statement,  "An  element  is  known  by  the 
company  it  keeps." 

22.  (a)  Compare  the  physical  properties  of  the  members  of 
the  halogen  family,    (b)  Compare  their  chemical  properties. 

23.  In  what  ways  do  the  halogens    (a)  resemble  one  an- 
other, and  (b)  differ  from  one  another? 

24.  List  the  halogens  in  the  order  of  their  chemical  activity. 


25.  Illustrate  what  is  meant  by  the  replacement  power  of 
the  halogens. 

26.  Does  the  following  equation  represent  an  actual  chem- 
ical reaction?  Explain.  2KC1  +  Br2 >2KBr  +  C12 

27.  Describe   two  uses  of  the  replacement  power  of  the 

28.  Make  a  labeled  diagram  of  the  laboratory  preparation 

of  Br2. 

29.  (a)  What  is  a  sedative?   (b)  State  three  other  uses  of 


30.  Describe  the  commercial  preparation  of  Br2. 


31.  Using  a  labeled  diagram,  describe  the  laboratory  prep- 
aration of  I2. 

32.  (a)  What    is    sublimation?     (b)  What    is    tincture    of 

33.  How  is  iodine  obtained  for  industrial  use? 

34.  How  is  the  most  active  of  all  the  chemical  elements 

35.  (a)  What  is  one  cause  of  mottled  teeth?    (b)  What  is 

36.  What  is  the  relationship  of  F2  to  dental  caries? 

37.  (a)  What  compounds  of  F2  are  superior  refrigerants? 
(b)  Why? 

Group  B 

38.  Compare  C12  and  O2   with  respect   to:     (a)  chemical 
activity,     (b)  behavior   with    H2,     (c)  valence.    Explain   each 
answer  fully. 

39.  (a)  What  is  the  relationship  between  lack  of  iodine 
and  goiter?   (b)  Why  is  goiter  not  so  prevalent  in  New  York 
City  as  it  is  in  some  other  parts  of  the  United  States? 

40.  A  bottle  of  tincture  of  iodine  was  found,  after  long 
use,  to  contain  only  a  dark  solid,   (a)  Would  it  be  safe  to  use 
it  after  adding  pure  ethyl  alcohol?   (b)  Explain  your  answer. 

41.  If  you  had  some  Nal,  how  would  you  prepare  a  solution 
of  tincture  of  iodine  from  it? 

42.  I2  is  produced  from  California  oil-well  brines  and  from 
other  brines  in  Michigan.  Can  you  suggest  a  way  in  which 
this  I9  is  extracted? 


43.  How  much  fluorine  would  be  needed  to  make  one  ton 
of  dichlorotetrafluorethane,  C2C12F4? 

44.  (a)  What  factors  must  a  manufacturer  consider  when 
he  chooses  raw  materials  for  use  in  manufacturing  a  substance 
on  a  large  scale?  (b)  Show  how  these  factors  apply  in  the  man- 
ufacture of  C12. 


1.  Make  a  study  of  advertisements  of  antiseptics  and  dis- 
infectants in  magazines  and  newspapers.  What  appeals  are 
made  in  these  advertisements  to  induce  you  to  buy  a  particular 
brand?  Are  the  appeals  chiefly  scientific,  pseudo-scientific,  fear- 
provoking,  or  do  they  appeal  chiefly  to  your  pride,  sense  of 
superiority,  your  desire  for  social  approval,  and  so  forth?  Il- 
lustrate your  report  with  actual  advertising  copy. 

2.  Ink  eradicators  frequently  contain  two  solutions.  No.  1 
contains  a  solution  of  a  very  weak  acid  and  No.  2  contains  a 
solution  of  sodium  hypochlorite.  Prepare  such  an  ink  eradi- 
cator  using  vinegar  or  citric  acid  for  No.  1  and  Javelle  water 
or  Clorox  for  No.  2.  Demonstrate  its  use  before  your  class. 

3.  Victor  Meyer,  an  eminent  German  chemist,  prepared  a 
compound  in  1886  which  is  now  known  as  mustard  gas.  Meyer 
was  the  first  chemist  to  prepare  this  chemical  during  his  the- 
oretical investigations.  Thirty  years  later  chemists,  cooperating 
with  the  Germany  military  machine,  made  this  compound 
available  for  use  as  a  poison  gas.  Should  society  take  a  hand  in 
suppressing  such  discoveries  which  might  be  used  against  man- 
kind? Write  a  brief  paper  either  in  favor  of  this  point  of  view 
or  against  it,  or  arrange  in  class  for  a  discussion  or  debate  on 
this  topic. 

4.  Study  the  quotation  at  the  beginning  of  this  chapter. 
Prepare  a  brief  report  to  the  class  on  the  meaning  of  this 
quotation.  Illustrate  your  report  with  an  example  from  the 
history  of  science. 



.  .  .  Our  experiences  and  observa- 
tions alone  never  lead  to  finalities. 
Theory,  however,  creates  reliable 
roads  over  which  we  may  pursue  our 
journeys  through  the  world  of  ob- 
servation. Anton  Reiser,  1930 

The  electron  theory  gives  us  a  clearer  picture  of  matter  and  its 
changes.  Formulating  accurate  theories  takes  remarkably  clear  in- 
sight, courage,  and  creative  imagination.  The  theories  and  principles 
of  science  are  among  the  most  truly  creative  products  of  the  mind  of 
man.  The  atomic  theory  of  Dalton  is  one  of  the  great  theories  upon 
which  modern  chemistry  is  built.  It  shows  us  that  atoms  do  not  com- 
bine in  a  haphazard,  irregular  manner  but  form  molecules  in  ac- 
cordance with  unvarying  natural  laws.  However,  it  gives  us  no  idea 
why  this  is  so. 

A  more  recently  developed  theory,  which  is  called  the  electron  the- 
ory, supplements  the  atomic  theory  and  gives  this  explanation.  It  pro- 
vides answers  also  to  such  questions  as  why  the  extreme  chemical 
activity  of  fluorine,  the  comparative  inactivity  of  nitrogen,  and  the 
inertness  of  argon;  why  elements  and  radicals  possess  the  valences 
that  they  have,  and  many  other  questions.  The  electron  theory  is  the 
fruit  of  many  scientists  who  worked  in  many  countries.  It  made  possi- 
ble the  atomic  age. 

Static  electricity.  About  2600  years  ago,  the  Greeks  discovered  that 
when  amber  is  rubbed  with  cloth,  it  becomes  capable  of  attracting 
tiny  bits  of  straw  or  dry  leaves.  Through  the  centuries,  men  have 



rubber  rod 

Fig.   26.    Demonstration    of   static   electricity. 

glass  rod 

discovered  that  other  materials  may  be  given  this  property.  Glass, 
when  rubbed  with  silk,  or  hard  rubber,  when  rubbed  with  fur,  will 
also  attract  light  objects.  The  force  which  causes  this  attraction  was 
named  electricity  from  the  Greek  word  for  amber.  Today,  we  refer 
to  the  electricity  caused  by  friction  or  rubbing  as  static  electricity. 

Benjamin  Franklin  attempts  to  explain  negative  and  positive 
electricity.  Suspend  a  pith  ball  by  a  silk  thread.  Touch  it  with  a  hard- 
rubber  rod  which  has  been  rubbed  with  fur.  As  soon  as  contact  is 
made,  the  pith  ball  will  be  driven  away,  or  repelled,  from  the  rod. 
Then  bring  near  the  pith  ball  a  glass  rod  which  has  been  rubbed 
with  silk.  The  ball  will  be  attracted  toward  the  glass  rod. 

This  simple  experiment  demonstrates  that  the  glass  and  the  hard- 
rubber  rods  were  charged  with  opposite  kinds  of  electricity.  When 
the  neutral  pith  ball  was  touched  by  the  hard-rubber  rod,  it  became 
charged  with  the  same  kind  of  electricity  as  the  rod.  It  was  then 
pushed  away  from  the  rod,  proving  that  objects  with  the  same  electric 
charges  repel  one  another.  However,  the  ball  was  drawn  toward  the 
glass  rod,  proving  that  objects  with  opposite  electric  charges  attract 
one  another. 

In  1747,  Benjamin  Franklin,  one  of  the  most  versatile  men  Amer- 
ica has  ever  produced,  received  a  static  electricity  machine  Irom  a 
friend  in  England.  Franklin,  son  of  a  soap-maker  who  had  fled  from 
England  because  of  religious  persecution,  was  then  41  years  old.  Be- 
cause of  a  very  successful  business  career,  he  was  rich  enough  to  re- 
tire from  business  and  devote  himself  to  scientific  experimentation. 
He  had  already  organized  the  first  scientific  society  in  the  New 
World,  an  organization  which  later  became  the  American  Philosophi- 
cal Society. 

Franklin  performed  many  experiments  with  his  static  electricity 
machine,  and  that  same  year  he  announced  his  own  views  on  the 
nature  of  electricity.  He  wrote:  "The  electric  fire  (electricity)  is  not 
created  by  friction,  but  collected,  being  really  an  element  diffused 
among  matter.  The  electrical  matter  consists  of  particles  extremely 
subtile.  .  .  .  Hence  have  arisen  some  new  terms  among  us:  we  say 
B  is  electrised  positively;  A,  negatively.  Or  rather  B  is  electrised  plus; 



A,  minus/'  Franklin  was  the  first  person  to  use  these  present-day 
terms  in  referring  to  electricity.  He  later  induced  his  good  friend 
Priestley  to  write  a  history  of  electricity,  and  thus,  in  part,  directed 
Priestley's  scientific  career. 

Franklin's  electric  theory  was  not  altogether  correct,  for  he  be- 
lieved that  if  a  body  has  too  much  electricity  it  is  charged  positively 
(-f )  ;  if  it  has  not  enough,  it  is  charged  negatively  (— );  and  if  it  has 
just  enough,  it  is  neutral.  Even  though  his  ideas  were  not  altogether 
correct,  his  reasoning  and  his  terminology  for  electricity  were  more 
modern  than  those  of  any  other  eighteenth-century  scientist.  So 
great  was  the  creative  imagination  of  Franklin  that  he  came  very 
close  to  arriving  at  the  modern  concept  of  the  electric  nature  of  mat- 
ter, a  point  of  view  reached  only  after  some  150  years  of  further  ex- 

Today  we  refer  to  the  charge  which  was  produced  on  the  hard- 
rubber  rod  mentioned  above  as  negative  (— ) ,  and  that  produced  on 
the  glass  rod  as  positive  (+) . 

The  electron  is  discovered.  Almost  a  century  and  a  half  after 
Franklin,  William  Crookes,  an  Englishman,  studied  the  effect  of 
passing  a  current  of  high-voltage  electricity  through  a  glass  tube  from 
which  nearly  all  the  air  had  been  pumped.  He  noticed  that  a  beam  of 
light  issued  from  the  negative  plate,  or  cathode,  of  the  tube.  When  a 
magnet  was  brought  near  the  tube,  the  beam  would  bend.  Since  ordi- 
nary light  is  not  affected  by  a  magnet,  the  beam  showed  a  property  of 
matter  rather  than  of  light. 

Another  Englishman,  J.  J.  Thomson,  undertook  to  explain  the 
strange  behavior  of  these  cathode  rays.  In  1897,  after  20  years  of  bril- 
liant research,  he  announced  his  results.  He  said  that  cathode  rays 
are  composed  of  particles  of  negative  electricity,  torn  away  from  the 
atoms  of  the  air  in  the  tube.  To  these  particles,  Thomson  gave  the 
name  electrons.  The  cathode  ray  was  bent  because  the  negative  elec- 
trons would  be  attracted  by  the  positive  pole  and  repelled  by  the 
negative  pole  of  a  magnet. 

Fig.  27.  Crookes'  tube.  Notice  how  the  positive  pole 
of  the  magnet  deflects  the  cathode  ray. 

stream  of  electrons 

Crookes'  tube 

Joseph  John  Thomson  was 
born  at  Manchester  in  1856. 
He  followed  Rayleigh  as  head 
of  the  Cavendish  Laboratory 
of  Experimental  Physics  at 
Cambridge  University. 

Brumt  font  tt  c, 

The  discovery  of  the  electron  also  explained  the  phenomenon  of 
static  electricity.  When  one  object  is  rubbed  with  another,  electrons 
are  transferred.  Thus,  when  glass  is  rubbed  with  silk,  the  glass  loses 
electrons  to  the  silk  and  is  left  with  a  positive  charge;  amber,  when 
rubbed  with  fur,  takes  electrons  from  the  fur  and  becomes  negatively 

Thomson's  discovery  of  the  electron  completely  upset  the  theory 
that  the  atom  is  the  smallest  unit  of  matter,  since  electrons  were 
found  in  all  atoms.  It  also  proved  that  the  idea  of  an  indivisible  atom 
was  inaccurate. 

Subsequently,  Robert  A.  Millikan,  an  eminent  American  scientist, 
succeeded  in  computing  the  mass  of  a  single  electron.  He  found  it 
to  be  about  y^Vr  °*  ^ie  mass  °*  one  hydrogen  atom.  (Mass  is  the 
amount  of  matter  that  a  substance  contains.  It  does  not  vary,  as  does 

lirown  Brothers 

Robert  Andrews  Millikan  (1868- 
1 953)  won  the  Nobel  physics  prize 
in  1923  for  his  work  in  isolating 
and  weighing  the  electron. 




electron  (outside  the  nucleus) 
proton  (in  the  nucleus) 

Fig.  28.  The  hydrogen  atom 

weight,  with  the  gravitational  pull  of  the  earth.  However,  weight  is 
dependent  upon  the  amount  of  matter  a  substance  contains.) 

The  proton  is  discovered.  After  Thomson  discovered  the  electron, 
one  of  his  students,  Ernest  Rutherford,  began  to  ponder  over  the  na- 
ture of  the  rest  of  the  atom.  An  atom  itself  is  electrically  neutral. 
Surely,  he  thought,  in  the  electrically  neutral  atom  there  must  be 
some  positive  electricity  to  counterbalance  the  negative  electron. 

After  much  research,  which  began  in  1911,  Rutherford  determined 
that  all  atoms  contain  one  or  more  particles  of  positive  electricity 
which  he  named  the  proton.  The  proton,  which  is  a  positively 
charged  atom  of  hydrogen,  is  1837  times  as  heavy  as  the  electron. 

The  structure  of  the  atom  according  to  the  electron  theory.  Atoms 
of  the  various  elements  contain  different  numbers  of  electrons  and 
protons.  In  general,  however,  the  arrangement  of  these  particles  fol- 
lows a  similar  pattern  in  all  elements.  Rutherford  gave  us  the  first 
picture  of  the  structure  of  the  atom.  It  resembles  our  own  solar  sys- 
tem with  its  sun  and  revolving  planets.  The  "sun"  of  the  atom  is 
called  the  nucleus.  It  is  composed  partly  of  protons  around  which, 
at  a  relatively  great  distance,  revolve  planetary  electrons.  A  diagram 
of  the  simplest  of  the  atoms,  that  of  hydrogen,  is  shown  above. 

It  was  Rutherford  who  found  that  electrons  and  protons  were  not 
evenly  distributed  in  the  atom  but  that  the  heavy  protons  were  all 
located  in  the  center.  He  shot  helium  atoms  (alpha  particles) 
through  a  cluster  of  nitrogen  atoms  and  photographed  the  results 
by  means  of  a  cloud  chamber  and  special  camera  devised  by 
C.  T.  R.  Wilson,  another  English  scientist.  He  found  that  only  in 
an  extremely  few  cases  the  path  or  fog  track  of  the  helium  bullet 
was  not  straight,  but  was  thrown  sharply  off  its  course.  On  the  basis 
of  the  volume  of  the  nitrogen  atom  and  the  ratio  of  straight  fog 
tracks  to  bent  ones,  he  calculated  that  the  protons  must  all  be  con- 
Fig.  29.  (left)  Apparatus  for  producing  the  cloud  effect. 
The  upper  chamber  is  filled  with  nitrogen,  (right)  The 
fog  track  of  an  alpha  particle. 

.water  vapor  condenses 


track  of  particle 
suddenly  released 


water. .  . 

is  compressed  . 


centrated  in  the  tiny  nucleus  of  the  atom.  The  diameter  of  the  nu- 
cleus is  about  10o*ooo  that  of  the  whole  atom.  Helium  bullets  pass- 
ing through  the  rest  of  the  atom  met  no  solid,  positively  charged 

Size  of  the  atom.  Most  of  us  are  baffled  in  attempting  to  imagine 
the  size  of  the  particles  within  the  atom  and  the  distances  between 
them.  The  diameter  of  a  hydrogen  atom  is  about  40,000,000  inch. 
If  an  atom  were  magnified  about  30  trillion  times,  its  diameter  would 
be  about  10  miles.  At  the  center  would  be  the  nucleus,  about  the  size 
of  a  tennis  ball.  The  electrons,  each  about  the  size  of  a  hazelnut, 
would  revolve  about  the  nucleus  in  orbits  in  somewhat  the  same  way 
as  the  planets  of  our  solar  system  revolve  about  the  sun. 

This  description  is  actually  an  oversimplification  of  the  structure 
of  the  atom.  However,  it  tells  us  an  important  fact.  Atoms,  which 
compose  every  element  and  compound,  are  largely  empty  space! 

The  neutron:  A  notable  scientific  prediction.  All  of  the  protons 
of  an  atom  are  located  in  the  nucleus.  However,  not  all  of  the  elec- 
trons of  an  atom  are  planetary;  some  electrons,  too,  are  found  within 
the  nucleus.  But  how  can  negatively  charged  electrons  and  positively 
charged  protons  exist  side  by  side  in  the  nucleus?  To  explain  this, 
Dr.  W.  D.  Harkins  predicted  the  existence  of  a  new  particle.  An 
electron  within  the  nucleus,  he  said,  does  not  exist  as  a  separate  par- 
ticle, but  is  combined  with  a  proton  forming  an  electrically  neutral 
particle  which  he  named  the  neutron.  Since  the  weight  of  the  elec- 
tron is  extremely  slight,  it  may  be  disregarded  in  figuring  the  weight 
of  the  neutron.  The  neutron  has  been  found  to  have  about  the  same 
mass  as  the  proton. 

In  1920,  the  existence  of  the  neutron  was  theoretically  established 
by  W.  D.  Harkins  of  the  University  of  Chicago.  Twelve  years  later, 
it  was  actually  discovered  by  James  Chadwick,  working  in  Ruther- 
ford's laboratory.  The  ultimate  verification  of  Harkins'  work  by 
Chadwick  demonstrates  the  value  of  pure  theory  and  creative  imag- 
ination in  science.  For  his  part  in  the  discovery,  Chadwick  was 
knighted  in  1944. 

The  electron  and  proton  are  so  close  together  in  the  neutron  that 
the  volume  of  the  neutron  is  millions  of  times  smaller  than  that  of 
any  atom.  It  has,  therefore,  an  extremely  high  density.  This  fact  has 
been  used  to  explain  the  extremely  high  densities  of  certain  stars. 
Since  the  neutron  is  electrically  neutral,  electric  forces  do  not  repel 
it.  Therefore,  the  neutron  has  great  penetrating  powers. 

We  may  say  that,  in  general,  all  matter  is  composed  of  three  kinds 
of  fundamental  electric  particles:  electrons,  protons,  and  \neutrons. 


Scientists  have  discovered  other  electric  particles  within  the  atom, 
as  shown  in  Table  8,  but  the  three  mentioned  here  are  considered 
the  most  important. 

How  the  structure  of  an  atom  may  be  represented  graphically. 
The  atom  of  each  element  contains  a  particular  number  of  electrons, 
protons,  and  neutrons  different  from  the  number  in  the  atom  of 
every  other  element.  Two  American  scientists,  Lewis  and  Langmuir, 
developed  a  theory  of  the  arrangement  of  planetary  electrons  which 
explains  why  each  element  has  different  chemical  properties. 

According  to  this  theory,  the  electrons  outside  the  nucleus  (plane- 
tary electrons)  arrange  themselves  in  successive  rings,  or  shells.  The 
first  ring  is  complete  when  it  contains  2  electrons;  the  second  ring 
is  complete  when  it  contains  8  electrons;  the  third  ring  is  complete 
when  it  contains  18  electrons;  the  fourth,  when  it  contains  either  18 
or  32  electrons.  However,  the  outermost  ring  never  contains  more 
than  eight  electrons.  According  to  Lewis  and  Langmuir,  an  atom 
with  30  planetary  electrons  will  have  its  first  ring  complete  with 
2  electrons,  second  ring  complete  with  8  electrons,  third  ring  com- 
plete with  18  electrons,  and  fourth  ring  incomplete  with  2  electrons. 
What  is  the  electron  structure  of  an  atom  with  48  planetary  electrons? 

The  number  of  protons  in  the  nucleus  of  the  atom  is  equal  to  the 
number  of  planetary  electrons.  This  equality  keeps  the  atom  elec- 
trically neutral.  Thus,  an  atom  with  30  planetary  electrons  will  have 
30  protons  in  its  nucleus. 

Of  all  the  elements,  hydrogen  has  the  simplest  atom.  Its  nucleus 
consists  of  one  proton.  Revolving  about  this  nucleus  is  one  planetary 
electron.  The  nucleus  of  the  helium  atom  contains  two  protons  (2+) 
and  two  neutrons  (2n)  ;  two  planetary  electrons  (2—)  revolve  about 
it.  Fig.  30  shows  how  we  may  graphically  represent  the  helium  atom 
and  also  the  chlorine  atom  (atomic  number  17)  . 

The  periodic  table  of  Mendeleyeff.  In  1869,  Mendeleyeff  (men'- 
de-la'ef) ,  a  Russian  chemist,  published  a  table  of  the  elements  ar- 
ranged in  order  according  to  their  increasing  atomic  weights.  He 
noticed,  that  when  arranged  in  this  manner,  the  elements  fell  into 

Fig.  30.  Diagrams  of  the  helium  and  chlorine  atoms. 

Helium  Chlorine 






































eight  distinct  groups.  Within  each  group,  the  elements  have  similar 
physical  and  chemical  properties. 

Let  us  examine  part  of  this  table  carefully.  Note  that  hydrogen  is 
omitted,  and  that  the  table  begins  with  helium,  the  element  with  the 
next  heaviest  atomic  weight.  Lithium  (Li) ,  with  an  atomic  weight 
of  6.940  follows,  and  so  on  through  fluorine  (19.0)  .  These  eight  dis- 
similar elements  comprise  one  series  or  period. 

The  element  following  fluorine  in  order  of  atomic  weight  is  neon 
(Ne) .  It  has  chemical  properties  similar  to  those  of  helium  (He)  and 
falls  directly  below  it  in  the  table.  Directly  below  helium  is  argon 
(A)  with  similar  properties.  Thus  helium,  neon,  argon,  and  certain 
succeeding  elements  comprise  Group  Zero.  This  group  of  elements 
is  known  as  the  family  of  inert  gases  (see  Chapter  7) .  The  elements 
in  Group  VII  are  known  as  the  halogen  family  (see  Chapter  10) . 
The  other  groups  are  also  made  up  of  closely  related  elements. 

Moseley  discovers  the  law  of  atomic  numbers.  Prior  to  1912,  the 
numerical  position  of  an  element  in  a  table  of  atomic  weight  was 
called  the  atomic  number  of  the  element.  It  occurred  to  Rutherford 
that  this  number  might  also  represent  the  number  of  protons  in  the 
nucleus  of  the  atom.  One  of  his  students,  Henry  G.  J.  Moseley,  un- 
dertook to  find  out  whether  Rutherford's  idea  was  valid. 

Moseley's  experiments  bore  out  Rutherford's  theory.  The  atom 
of  each  element  was  found  to  contain  a  number  of  protons  in  its  nu- 
cleus corresponding  to  the  element's  numerical  position  in  the  peri- 
odic table  of  atomic  weights.  The  hydrogen  atom,  which  appears  first 
in  the  periodic  table,  has  only  one  proton  in  the  nucleus  of  its  atom; 
uranium,  which  appears  in  the  ninety-second  position  in  the  periodic 
table,  has  92  protons  in  the  nucleus  of  each  atom.  Moseley  showed 
that  the  atomic  number  of  any  element  is  equal  to  the  number  of  free 
protons  in  the  nucleus  of  its  atom.  This  is  known  as  the  law  of  atomic 
numbers.  Since  the  number  of  free  protons  in  the  nucleus  is  equal 
to  the  number  of  electrons  around  the  nucleus,  the  atomic  number 
of  an  element  is  also  equal  to  the  number  of  planetary  electrons. 

For  the  first  17  elements  in  the  table  of  atomic  numbers,  it  is  help- 
ful to  remember  that  the  atomic  number  is  equal  to  half  the  atomic 






weight  (disregarding  fractions).  Thus  the  atomic  number  of  chlo- 
rine (atomic  weight  35.457)  is  17. 

The  periodic  table  of  atomic  numbers.  Mendeleyeff's  periodic 
table  based  upon  atomic  weights  served  science  for  50  years.  In  1912, 
however,  it  was  displaced  by  the  new  periodic  table  developed  by 
Moseley  from  his  law  of  atomic  numbers. 

Moseley's  table  is  more  fundamental  than  MendeleyefFs  and  easily 
accounts  for  some  of  the  discrepancies  in  the  latter.  For  example,  the 
element  argon  has  an  atomic  weight  of  39.944  and  the  element  potas- 
sium an  atomic  weight  of  39.10.  Argon  should,  therefore,  follow 
potassium  in  the  table  based  on  atomic  weights.  But  the  properties 
of  argon  put  it  in  the  group  of  inert  elements,  preceding  potassium. 
Moseley's  researches,  which  showed  that  the  atomic  number  of  argon 
is  18  and  that  of  potassium  is  19,  eliminated  this  problem.  When 
Moseley  developed  his  table,  all  of  the  elements  had  not  been  discov- 
ered. Therefore  the  atomic  numbers  did  not  run  entirely  consecu- 
tively from  one  to  92.  Since  then  the  missing  elements  have  been  dis- 
covered, as  you  may  see  from  the  modern  periodic  table  below. 

A  new  definition  of  atomic  weight.  As  we  have  learned,  the  elec- 
tron, for  most  purposes,  may  be  considered  weightless,  and  the  proton 
and  neutron  may  be  considered  equal  in  mass.  Thus,  the  most  abun- 
dant hydrogen  atom,  which  contains  a  single  proton  and  no  neutrons, 
may  be  considered  equal  to  one  proton  in  weight;  the  helium  atom, 
which  contains  two  protons  and  two  neutrons,  weighs  four  times  as 
much  as  the  hydrogen  atom.  We  may  say  that  the  atomic  iveight  of 
hydrogen  is  one  and  the  atomic  weight  of  helium  is  four.  Thus,  ac- 
cording to  the  electron  theory,  the  atomic  weight  for  each  element 
may  be  defined  as  the  sum  of  the  protons  and  neutrons  in  the  nucleus 
of  an  atom  of  that  element.  Look  at  the  diagrams  of  various  atoms  in 
Fig.  32.  What  is  the  atomic  number  of  each?  the  atomic  weight? 

How  the  electron  theory  explains  isotopes.  In  1815,  William 
Prout,  a  London  physician,  announced  the  theory  that  all  the  chem- 
ical elements  are  made  up  of  groups  of  hydrogen  atoms  only.  Prout's 
theory  was  not  taken  seriously  for  100  years  until  Moseley's  work 
on  atomic  numbers  made  Prout's  conclusion  more  plausible. 

Since  the  nucleus  of  an  atom  of  any  element  is  composed  of  only 


'Following  lanthanum  are  1 4  elements  known  as  the  rare  earth  elements  (at.  no.  58-71). 
*  *  Following  uranium  are  eight  newly  created  elements  of  the  actinide  series  (at.  no.  93-100). 









52.01  34 



50.95  33 









95.95  52 






180.88  83 


183.92          84 


186.31          85 








Henry  6.  J.  Moseley  (1887- 
1915),  a  pupil  of  Ernest 
Rutherford,  discovered  the 
law  of  atomic  numbers  in 
1912.  His  brilliant  career  was 
ended  by  his  death  at  27 
during  World  War  I. 

neutrons  and  protons,  and  since  the  weight  of  each  of  these  units 
is  really  the  weight  of  the  hydrogen  atom,  it  may  have  occurred  to 
you  that  the  atomic  weights  of  all  the  elements  ought  to  be  whole 
numbers.  But  the  fact  that  many  atomic  weights,  for  example,  chlo- 
rine (35.457)  ,  are  not  whole  numbers  could  riot  be  brought  into 
harmony  with  this  idea. 

In  1913,  Theodore  W.  Richards  found  two  different  kinds  of  lead 
with  atomic  weights  of  206  and  207,  respectively,  and,  in  the  same 
year,  two  kinds  of  neon  with  different  atomic  weights  were  reported 
also.  The  name  isotopes  was  given  to  atoms  of  the  same  element  hav- 
ing the  same  chemical  properties  but  different  atomic  weights.  Dis- 
coveries of  isotopes  of  many  other  elements  soon  followed,  one  of 
which  (tin)  is  now  known  to  have  as  many  as  10  stable  isotopes  and, 
hence,  10  different  atomic  weights. 

The  discovery  of  isotopes  removed  the  obstacles  to  the  acceptance 
of  Prout's  idea.  For  elements,  as  we  know  them,  are  really  mixtures 
of  isotopes  having  different  atomic  weights,  each  of  which  is  a  whole 
number.  Thus  ordinary  chlorine  gas  is  really  made  up  of  some  atoms 
with  an  atomic  weight  of  35,  other  atoms  of  atomic  weight  37,  and 
still  others  of  atomic  weight  39.  Its  accepted  atomic  weight,  35.457, 
is  the  average  of  the  atomic  weights  of  the  three  different  weights  of 
chlorine  atoms  in  any  sample  of  the  gas.  Here  was  another  startling 
discovery  which  helped  to  destroy  Dalton's  idea  of  an  atom  whose 
atomic  weight  never  changed. 

The  electron  theory  explains  isotopes  as  caused  by  a  different  num- 
ber of  neutrons  in  each  kind  of  atom.  Thus,  isotopes  of  chlorine  with 
atomic  weights  of  35,  37,  and  39  behave  alike  chemically  because 
they  have  the  same  arrangement  of  planetary  electrons.  They  differ 

Fig.  31.  Diagrams  of  the  isotopes  of  chlorine. 


35     ' 










in  weight  because  of  a  difference  in  the  number  of  neutrons  in  their 
nuclei.  We  must,  then,  redefine  the  term  element.  A  substance  of 
which  all  atoms  have  the  same  atomic  number  is  an  element. 

How  the  electron  theory  explains  valence.  From  the  diagram  of  the 
chlorine  atom  (Fig.  31) ,  you  can  see  that  its  outermost  ring  contains 
7  electrons.  One  more  electron  is  needed  to  make  the  8  electrons 
needed  to  complete  this  ring. 

An  atom  whose  outermost  ring  is  nearly  complete  has  a  tendency 
to  borrow  enough  electrons  to  complete  this  ring.  An  atom  whose 
outermost  ring  has  few  electrons  tends  to  lose  electrons.  The  number 
of  electrons  gained  or  lost  by  an  atom  of  an  element  is  the  valence  of 
that  element.  Since  the  chlorine  atom  needs  to  borrow  only  1  elec- 
tron to  complete  its  outer  ring,  its  valence  is  1.  In  borrowing  this 
electron  it  becomes  negatively  charged  and,  as  a  result,  the  valence 
of  chlorine  is  negative.  Hence,  the  valence  of  chlorine  is  —1. 

An  atom  that  lends  electrons  becomes  positively  charged.  Hence, 
elements  whose  atoms  lend  electrons  have  positive  valences.  Thus, 
the  sodium  atom  with  an  atomic  number  of  11  (roughly  half  of 
22.997)  may  be  pictured  as  shown  in  Fig.  32.  As  you  see,  the  outer- 
most ring  contains  1  electron  which  the  atom  may  lend.  Hence  the 
valence  of  sodium  is  -f-1. 

An  atom  whose  outer  ring  is  complete  will  neither  lend  nor  bor- 
row electrons.  Elements  whose  atoms  are  of  this  type  have  a  valence 
of  0.  The  atom  of  neon  is  shown  in  Fig.  32. 

Flectrons  in  the  outermost  ring  of  an  atom,  which  may  be  either 
borrowed  or  lent,  are  called  valence  electrons. 

How  the  electron  theory  explains  metals  and  nonmetals.  A  metal 
is  a  lender  of  electrons.  That  is,  the  outermost  ring  of  an  atom  of  a 
metal  has  less  than  four,  or  half  the  number  (eight)  required  to 
complete  it.  When  such  an  atom  lends  electrons,  it  necessarily  be- 
comes positively  charged.  The  valence  of  metals,  therefore,  is  con- 
sidered positive. 

The  atom  of  a  nonmetal  is  a  borrower  of  electrons.  That  is,  its 
outermost  ring  has  more  than  four,  or  half  the  number  of  electrons 
(eight)  required  to  complete  it.  By  borrowing  electrons,  such  an 
atom  becomes  negatively  charged.  The  valence  of  nonmetals  is, 
therefore,  negative. 

Fig.  32.  Diagrams  of  various  atoms. 







If  the  outermost  ring  of  an  atom  of  an  element  has  just  half  the 
number  of  electrons  required  to  complete  it,  it  may  either  borrow  or 
lend  electrons.  Such  an  element  is  said  to  be  amphoteric.  A  common 
example  of  an  amphoteric  element  is  carbon  (atomic  weight  12) , 
whose  diagram  appears  in  Fig.  32. 

How  the  electron  theory  explains  electric  currents.  According  to 
'modern  theory,  an  electric  current  is  a  flow  of  electrons.  Atoms  of 
metals,  such  as  copper,  silver,  and  gold,  are  good  conductors  of  elec- 
tricity, because  some  of  their  electrons  are  held  loosely,  and  can  move 
freely  through  the  solid.  In  general,  nonmetals  are  poor  conductors, 
because  their  electrons  are  not  held  as  loosely  as  those  of  metals. 

How  the  electron  theory  explains  chemical  activity.  An  atom 
tends  to  complete  its  outer  ring  of  electrons.  If  an  element  such  as 
neon  or  argon  already  has  its  outer  ring  complete,  that  element  is 
inert.  That  is,  its  atom  will  not  lend  or  borrow  electrons,  and  hence 
the  element  is  completely  inactive  chemically.  In  general,  the  smaller 
the  number  of  electrons  an  atom  must  either  borrow  or  lend  to  com- 
plete its  outer  ring  of  electrons.,  the  greater  is  the  chemical  activity 
of  that  atom. 

An  atom,  then,  with  either  1  or  7  electrons  in  its  outermost  ring 
is  extremely  active.  Such  an  atom  is  fluorine,  whose  atomic  weight 
is  19,  and  atomic  number  is  9.  It  has  7  electrons  in  its  second  ring 
and  will  borrow  1  more  electron.  Potassium,  whose  atomic  weight 
is  39,  has  only  1  electron  in  its  fourth  ring,  and  hence  can  lend  only 
1  electron.  Atoms  such  as  those  ot  oxygen  and  magnesium  have  2 
electrons  to  borrow  or  lend  and  are  quite  active.  Atoms  of  nitrogen 
and  aluminum  have  3  electrons  to  borrow  or  lend  and  are  not  very 
active.  Generally,  the  farther  away  the  outer  ring  of  an  atom  is  from 
the  nucleus,  the  less  is  the  attraction  of  the  nucleus  for  its  electrons. 

This  helps  explain  the  chemical  behavior  ot  metals.  For  exam- 
ple, potassium  is  more  active  than  sodium  since  the  electron  in  its 
fourth  ring  can  be  lost  more  easily  than  the  electron  in  the  third 
ring  of  sodium.  Conversely,  the  closer  the  outer  ring  of  an  atom  is 
to  the  nucleus,  the  stronger  is  the  attraction  of  the  nucleus  for  its 
electrons.  This  fact  helps  explain  the  chemical  behavior  of  non- 
metals.  Fluorine  is  more  active  than  iodine  since  its  second  ring  has 

Fig.  33.  Chemical  union  of  sodium  and  chlo- 
rine according  to  the  electron  theory. 



sodium  chloride 


a  greater  attraction  for  the  electrons  of  other  atoms  than  does  the  fifth 
ring  of  iodine.  Both  of  these  rules  should  be  considered  rough  guides. 
Many  exceptions  occur  since  the  whole  problem  of  chemical  activity 
is  quite  complex. 


1  .  The  at.  wt.  of  sulfur  is  32.  Make  a  picture  of  its  atom  ac- 
cording to  the  electron  theory,  and  explain  its  valence  and 
chemical  activity. 

2.  With  the  aid  of  a  diagram,  show  why  helium,  at.  wt.  4, 
is  inert. 

3.  With  the  aid  ot  diagrams,  show  why  lithium,  sodium, 
and  potassium  belong  to  the  same  family  ot  elements. 

4.  What  are  the  valence  and  chemical  activity  of  an  element 
whose  outer  ring  contains  4  electrons? 

5.  With  a  diagram,  explain  the  valence  of  the  (OH)  radical. 

How  the  electron  theory  explains  chemical  union  and  electro- 
valence.  Chemical  activity  is  the  tendency  ot  atoms  to  complete  their 
outer  rings  and  form  stable  compounds.  Chemical  union  is,  there- 
fore, the  shifting  or  sharing  of  electrons  in  the  outer  electron  rings 
until  a  stable  condition  is  reached. 

A  metallic  atom  with  a  valence  of  +  1  exhibits  a  strong  attraction 
tor  a  non  metal  lie  atom  whose  outer  ring  needs  1  electron  to  com- 
plete it.  For  example,  the  union  of  sodium,  a  very  active  metal  with 
a  valence  ot  +1,  with  chlorine,  a  very  active  nonmetal  with  a  valence 
of  —  1,  may  be  represented  thus: 

The  extra  electron  on  the  outer  ring  of  the  sodium  atom  shifts 
over  to  the  vacant  space  in  the  almost  complete  outer  ring  of  the 
chlorine  atom.  Now  the  two  outer  rings  are  both  complete  and  the 
resulting  compound,  sodium  chloride,  is  very  stable.  Its  water  solu- 
tion conducts  electricity.  Compounds  formed  by  the  shifting  of  single 
electrons  are  polar  or  ionic  compounds.  Their  valence  is  called  elec- 

Fig.   34.   Formation   of  a   molecule   of  fluorine 
according  to  the  electron  theory. 

shared  pair  of  electrons 
jf  • 


How  the  electron  theory  explains  covalent  compounds.  In  many 
chemical  reactions,  there  is  no  actual  shifting  of  single  electrons  but 
rather  an  equal  sharing  of  a  pair  or  pairs  of  electrons.  The  com- 
pounds formed  are  nonpolar  compounds  and  the  valence  is  called 
covalence.  This  kind  of  combination  is  generally  very  strong  and  the 
molecules  so  formed  hold  together  well.  Nonpolar  compounds  are 
stable  and  generally  do  not  conduct  electricity.  Many  organic  com- 
pounds, such  as  alcohol,  glycerin,  and  sugar,  are  nonpolar. 

The  fluorine  molecule  (see  Fig.  34)  illustrates  the  sharing  of  a 
pair  of  electrons.  There  is  neither  a  gain  nor  a  loss  of  electrons  — 
simply  a  sharing.  Molecules  of  other  gases  that  consist  of  2  atoms, 
such  as  oxygen  and  chlorine,  exhibit  this  same  sharing  of  electrons. 
Polar  and  nonpolar  compounds  are  discussed  more  fully  on  page  238. 

How  the  electron  theory  explains  oxidation  and  reduction.  In 
the  equation  CuO  +  H2  -^  Cu  +  H2O,  copper  oxide  is  reduced  to 
copper  and  hydrogen  is  oxidized  to  water.  The  valence  of  Cu  has 
changed  from  plus  two  (in  CuO)  to  zero  in  free  copper,  and  the  cop- 
per has  gained  2  electrons  (Cu++  +  2e  -»  Cu°) .  The  hydrogen  has 
lost  an  electron  and  changed  to  H+  in  H2+O —  (H°  —  e  -»  H+) . 

Oxidation  has  been  defined  thus  far  as  the  union  of  a  substance 
with  oxygen.  Reduction,  similarly,  has  been  defined  as  the  removal 
of  oxygen  from  a  compound.  From  the  viewpoint  pf  the  electron 
theory  these  terms  take  on  much  broader  meanings.  A  loss  of  elec- 
trons resulting  in  an  increase  in  the  valence  of  an  element  is  called 
oxidation.  A  gain  of  electrons  resulting  in  a  decrease  in  the  valence 
of  an  element  is  called  reduction. 

Oxidation  and  reduction  in  this  broader  sense  need  not  involve 
either  oxygen  or  hydrogen.  Thus,  in  the  replacement  of  the  iodine 
in  potassium  iodide  by  chlorine,  as  explained  in  the  preceding  chap- 
ter, we  have  an  oxidation-reduction  reaction.  According  to  the  elec- 
tron theory,  this  reaction  is: 

K+I-  +  Cl°  ->  K+C1-  +  1° 

The  iodine  in  KI  has  lost  an  electron  and  changed  to  free  iodine,  1°. 
It  has  been  oxidized  by  chlorine  which  has  gained  an  electron  and 
changed  from  free  chlorine  to  the  negative  chloride  ion.,  Cl~  (see 
page  232) . 

How  the  electron  theory  aids  in  balancing  equations.  We  may  see 
from  this  oxidation-reduction  equation  how  each  atom  of  iodine  lost 
an  electron  and  was  oxidized  and  how  each  free  chlorine  atom  gained 
an  electron,  forming  a  chloride  ion  with  a  negative  charge.  This  was 
a  relatively  simple  reaction.  But  how  is  a  more  complex  oxidation- 


reduction  equation  balanced  and  how  does  our  understanding  of  the 
electron  theory  help  us  to  find  correct  coefficients  to  use  in  balanc- 
ing the  equation?  Let  us  consider  some  examples: 

EXAMPLE  A:  It  is  desired  to  reduce  ferric  chloride,  FeCl3,  to 
ferrous  chloride,  Fed,,  by  the  use  of  the  evil-smelling  gas, 
H2S  —  a  common  reducing  agent.  In  the  reaction,  free  sulfur 
and  hydrochloric  acid  are  formed  also. 

1)  Write  the  unbalanced  equation  for  the  reaction. 

FeCl8  +  H2S  -»  FeCl2  +  S  +  HC1 

2)  Select  the  atoms  which,  according  to  the  electron  theory, 
are  either  reduced  or  oxidized,  that  is,  either  gain  or  lose  elec- 
trons. Place  the  valences  involved  at  the  upper  right  of  the 
symbol  of  each  atom. 

Fe+++Cl3  +  H2S—  ->  Fe++Cl2  +  S°  +  HC1 

3)  By  inspection,  we  see  that  each  atom  of  iron  gains  one 
electron  and  is  reduced;  each  atom  of  sulfur  loses  two  electrons 
and  is  oxidized.  The  number  of  electrons  involved  in  the  oxi- 
dation of  H2S  is  twice  as  great  as  the  number  of  electrons  in- 
volved in  the  reduction  of  one  molecule  of  FeCl3.  Therefore, 
twice  as  many  molecules  of  FeCl3  must  have  been  reduced. 
Indicate  this  by  writing  the  coefficient  2  before  FeCl3  in  the 
equation.  Then  balance  the  equation  by  the  methods  you  have 
already  learned. 

2FeCl8  +  H2S  -»  2FeCl2  +  S  +  2HC1 

In  actual  practice,  chemists  usually  consider  only  the  changes  in 
valence  of  the  atoms  involved  rather  than  the  shifting  or  sharing  of 
electrons.  The  sum  of  the  changes  in  valence  on  both  sides  of  the 
arrow  must  be  equal.  Consequently,  this  method  of  balancing  equa- 
tions is  known  as  the  valence-change  method. 

EXAMPLE  B:  Using  the  valence-change  method,  write  the  equa- 
tion for  the  reaction  between  the  oxidizing  agent,  potassium 
permanganate,  and  hydrochloric  acid.  The  products  are  potas- 
sium chloride,  manganese  chloride,  water,  and  free  chlorine. 

1)  Write  the  unbalanced  equation,  omitting  the  subscript 
from  the  free  chlorine.  It  is  not  necessary  to  indicate  the 
valence  of  those  atoms  which  retain  the  same  valence  through- 
out the  reaction. 

4  +  HC1-  -»  KC1  4-  Mn++Cl2  +  H2O  +  Cl° 


2)  Note  that  the  following  changes  occur: 

a)  Each  atom  of  manganese  in  KMnO4  gains  five  elec- 
trons in  combining  with  chlorine  to  form  MnCl2;  therefore  the 
valence  loss  of  the  manganese  is  five. 

(b)  Each  atom  of  chlorine  loses  one  electron  in  becoming 
free  chlorine;  therefore  the  valence  gain  of  each  atom  of  lib- 
erated chlorine  is  one. 

3)  The  five  valences  lost  by  the  manganese  must  be  balanced 
by  the  valence  gained  by  the  chlorine  since  no  other  element 
changes  valence  in  the  reaction.  To  do  this,  we  must  show  that 
for  each  molecule  ot  KMnO4  used,  five  atoms  of  Cl  are  lib- 
erated. Therefore,  write  the  coefficient  5  before  the  free  Cl: 

KMn04  +  HC1  ->  KC1  +  MnCl2  +  H2O  +  5C1 

4)  As  we  know,  free  chlorine  exists  as  a  molecule  composed 
of  two  atoms,  therefore  we  add  the  subscript,  2,  to  Cl.  In  order 
to  keep  the  valence  changes  in  balance,  we  consider  that  two 
molecules  of  KMnO4  are  used  in  liberating  five  molecules  of 
C12.  Therefore,  write  the  coefficient  2  before  KMnO4. 

2KMn04  +  HC1  ->  KC1  +  MnCl*  +  H2O  +  5C12  f 

5)  Complete  the  balancing  of  the  equation  in  the  usual 
manner.  Check  your  work  carefully. 

2KMnO4  +  16HC1  ->  2KC1  +  2MnCl2  +  8H2O  +  5C12  T 

From  this  discussion,  we  can  formulate  the  following  method  of 
balancing  oxidation-reduction  equations: 

1)  Write  the  unbalanced  equation. 

2)  Write  the  changes  in  valence  at  the  upper  right  of  the 
symbols  of  the  atoms  that  are  oxidized  and  those  that  are  re- 

3)  Find  the  valence  gains  and  losses  and  decide  on  the  coef- 
ficients that  will  make  them  equal.  Remember  that  molecules 
composed  of  two  atoms  may  be  involved  and  adjust  the  coeffi- 
cients accordingly. 

4)  Complete  the  balancing  of  the  equation  by  the  usual 

Changing  theories  and  the  spirit  of  modern  science.  In  1924,  the 
electron  theory,  backed  by  a  mass  of  experimental  evidence,  was 
quite  generally  accepted.  In  that  year,  Prince  de  Broglie  (de  bro'y') 
suggested  that  the  electron  is  not  merely  a  particle  of  electricity,  as 
the  electron  theory  explained,  but  like  light,  is  composed  of,  possesses, 
or  perhaps  is  attended  by,  a  group  of  waves. 


By  1927,  two  Americans,  Davisson  and  Germer,  proved  de 
Broglie's  theory  experimentally,  showing  that  both  electrons  and 
protons  possess  a  property  of  light  (a  wave  phenomenon) .  In  1929, 
de  Broglie  was  honored  with  the  Nobel  prize  in  physics  for  his  theory. 

The  theory  that  matter  possesses  wave  properties  created  an  up- 
heaval in  the  existing  theories.  However,  we  still  find  it  convenient 
to  regard  the  electron,  proton,  and  neutron  as  tiny  individual  par- 
ticles. If  new  facts  show  we  are  wrong,  we  shall  scrap  this  concept. 

In  every  phase  of  science,  this  practice  is  followed.  Old  ideas  are 
retained  as  long  as  they  are  useful.  They  may  be  altered  somewhat 
to  fit  newly-discovered  facts.  However,  if  enough  new  data  are  accu- 
mulated to  prove  them  incorrect,  the  old  theories  are  abandoned. 
It  is  the  belief  of  scientists  that  tradition  should  never  be  allowed  to 
stand  in  the  way  of  greater  enlightenment. 

Albert  Einstein,  one  of  the  most  eminent  of  living  scientists,  when 
he  was  suddenly  confronted  with  new  facts  which  could  not  fit  his 
own  theories,  expressed  the  spirit  of  science  thus:  "The  new  facts 
have  smashed  my  old  ideas  like  a  hammer  blowl"  And  he  went  on 
to  change  some  of  his  most  cherished  theories. 


Jaffe,  Bernard.  Crucibles:  The  Story  of  Chemistry,  pp.  289- 
311.  Simon  and  Schuster,  New  York,  1948.  A  sketch  of  Henry 
Moseley,  whose  lifework  was  done  in  4  short  years.  Before  the 
world  knew  this  genius,  he  died. 

Moulton,  Forest  R.,  and  Schifferes,  Justus  J.,  Ed.  The  Auto- 
biography of  Science,  pp.  502-509.  Doubleday  Doran  &  Co., 
New  York,  1945.  Contains  the  original  papers  dealing  with  the 
origin  of  the  electron  theory. 


1 .  Electric  charges  caused  by  rubbing  or  friction  are  called 
static  electricity. 

2.  Electrons  are  tiny  particles  of  negative  electricity.  A 
cathode  ray  is  a  stream  of  electrons.  Our  knowledge  of  electrons 
is  the  result  of  the  work  of  William  Crookes,  J.  J.  Thomson, 
Robert  A.  Millikan,  and  others. 

3.  Protons  are  particles  oi  positive  electricity.  A  proton  is 
1837  times  heavier  than  an  electron.  It  was  discovered  and 
named  by  Ernest  Rutherford. 


4.  A  neutron  is  an  electrically  neutral  particle  composed  of 
one  electron  and  one  proton.  Its  mass  is  about  the  same  as 
that  of  a  proton,  or  of  the  atom  of  the  lightest  hydrogen 

5.  All  atoms  are  composed  of  electrons,  protons,  and  neu- 
trons with  the  exception  of  the  atom  of  the  lightest  hydrogen 
isotope  which  is  composed  of  one  electron  and  one  proton. 
Protons  and  neutrons  are  located  in  the  nucleus  of  the  atom. 
Electrons  revolve  about  the  nucleus  in  orbits.  The  relative 
distance  between  the  nucleus  and  the  planetary  electrons  is 
so  great  that  we  may  say  the  atoms  of  all  elements  are  largely 
empty  space. 

6.  A  fixed  number  of  electrons  is  required  to  complete 
each  electronic  orbit  of  the  atom.  The  first  ring  is  complete 
when  it  contains  2  electrons;  the  second,  when  it  contains  8; 
the  third,  when  it  contains  18;  the  fourth,  when  it  contains  32. 
The  outermost  ring  never  contains  more  than  8  electrons. 

7.  The  atomic  number  of  any  element  is  the  number  of 
free  protons  in  the  nucleus  of  its  atom.  According  to  the  law 
of  atomic  numbers,  the  elements  can  be  arranged  in  a  periodic 
table  in  the  order  of  their  increasing  atomic  numbers. 

8.  The  atomic  weight  of  any  element  is  the  sum  of  the 
protons  and  neutrons  in  the  nucleus  of  one  of  its  atoms. 

9.  Isotopes  are  different  forms  of  the  s^me  element.  They 
possess  the  same  chemical  properties,  but  have  slightly  different 
physical  properties.  Isotopes  of  an  element  all  have  the  same 
atomic  number,  but  different  atomic  weights. 

10.  The  valence  of  an  element  is  the  number  of  electrons 
that  its  atom  must  borrow  or  lend  to  complete  its  outermost 
ring.  Electrons  in  the  outermost  ring  of  an  atom  which  may 
be  borrowed  or  lent  are  valence  electrons. 

11.  An  element  is  metallic  if  the  outer  ring  of  its  atom  con- 
tains less  than  half  the  number  of  electrons  necessary  to  com- 
plete the  ring.  In  other  words,  metals  are  lenders  of  electrons. 

12.  An  element  is  nonmetallic  if  the  outer  ring  of  its  atom 
contains  more  than  half  the  number  of  electrons  necessary  to 
complete  the  ring.  In  other  words,  nonmetals  are  borrowers 
of  electrons. 

13.  An  element  is  amphoteric  if  the  outer  ring  of  its  atom 
contains  exactly  half  the  number  of  electrons  necessary  to  com- 
plete this  outer  ring. 

14.  An  electric  current  is  a  flow  of  electrons. 

15.  Chemical  activity  depends  upon  the  number  of  elec- 
trons in  the  outermost  ring  of  the  atom  of  an  element.  When 
the  outermost  ring  is  complete,  the  element  is  inert;  if  the 
outermost  ring  lends  or  borrows  one  electron,  the  element  is 


very  active;  if  the  outermost  ring  lends  or  borrows  three  elec- 
trons, the  element  is  not  very  active. 

16.  Chemical  union  is  the  shifting  or  sharing  of  the  electrons 
in  the  outer  rings  until  a  stable  condition  is  reached. 

17.  Electrovalent  compounds  are  those  formed  by  the  shift- 
ing of  single  electrons;  covalent  compounds  are  those  formed 
by  the  sharing  of  a  pair  or  pairs  of  electrons. 

18.  Oxidation,  a  loss  of  electrons,  increases  the  valence  of 
an  element.   Reduction,   a   gain   of  electrons,  decreases   the 
valence  of  an  element. 

19.  New  theories  are  constantly  being  advanced  about  the 
nature  of  matter.  Recently  it  has  been  suggested  that  both  the 
proton  and  electron  possess  wave  properties.  It  is  necessary  for 
scientists  to  be  ready  to  give  up  theories  when  facts  show  that 
newer  theories  are  more  nearly  accurate. 


Group  A 

1.  (a)  How  may  static  electricity  be  produced?    (b)  How 
did  the  discovery  of  electrons  explain  this  phenomenon? 

2.  (a)  Who  discovered  the  electron?  (b)  How  did  his  dis- 
covery explain  the  behavior  of  cathode  rays?   (c)  Who  deter- 
mined the  mass  of  the  electron?   (d)  What  were  his  findings? 

3.  (a)  Who  discovered   the  proton?    (b)  What  principle 
of  electricity  led  him  to  his  discovery?   (c)  Do  the  protons  of 
different  elements  differ? 

4.  (a)  What  is  the  neutron?  (b)  What  scientists  took  part 
in  its  discovery?   (c)  Why  is  the  neutron  considered  approxi- 
mately equal  in  mass  to  the  proton? 

5.  (a)  Describe  the  general  arrangement  of  protons,  elec- 
trons, and  neutrons  in  an  electrically  neutral  atom,    (b)  De- 
scribe the  atom  in  terms  of  the  solar  system. 

t .  .  . 

6.  Explain  how  elements  differ  according  to  the  electron 

7.  (a)  Who  developed  the  first  periodic  table  of  atomic 
weights?   (b)  What  was  learned  from  arranging  the  elements 
in  such  a  table? 

8.  (a)  What  relationship  between  atomic  number  and  pro- 
tons was  demonstrated   by   Moseley?    (b)  Why   is   Moseley's 
periodic  table  more  fundamental  than  Mendeleyeff  s? 


9.  (a)  For  the  first  17  elements,  what  is  the  general  re- 
lationship between  atomic  number  and  atomic  weight?  (b)  Il- 
lustrate your  answer. 

10.  (a)  Define  atomic  weight  in  terms  of  the  electron  the- 
ory, (b)  Why  do  most  elements  have  atomic  weights  which  are 
not  whole  numbers? 

11.  (a)  What  is  an  isotope?    (b)  Explain  why  all  isotopes 
of  an  element  have  the  same  chemical  properties. 

12.  (a)  How    does    the    electron    theory    explain    valence? 
(b)  How  does  it  explain  the  behavior  of  metals?   (c)  of  non- 
metals?   (d)  of  amphoteric  elements? 

13.  Explain  the  relationship  between  the  number  of  elec- 
trons in  the  outer  ring  of  an  atom  and  the  chemical  activity 
of  that  element. 

14.  (a)  Make  a  diagram  of  the  structure  of  the  element 
whose  atomic  number  is  20.  (b)  Describe  some  of  the  element's 
chemical  properties. 

15.  Phosphorus  (P)  has  an  atomic  weight  of  30.98.  (a)  Dia- 
gram  the   structure  of  its   atom,    (b)  Describe   its   chemical 

16.  An  atom  has  a  nucleus  containing  1?  protons  and  18 
neutrons,  (a)  Make  a  structural  diagram  of  the  atom,  (b)  De- 
scribe its  chemical  properties. 

17.  The  atomic  weight  ot  beryllium  (Be)  is  9.02.   (a)  What 
are  the  atomic  number,  valence,  and  chemical  properties  of 
beryllium?  (b)  Is  it  a  metal  or  nonmetal? 

18.  The  atomic  weight  of  curium,  one  of  the  newly-discov- 
ered elements,  is  242;  its  atomic  number  is  96.  How  many 
neutrons  are  in  the  nucleus  of  one  of  its  atoms? 

19.  Make  a  diagram  of   (a)  the  oxygen  atom  and    (b)  the 
sulfur  atom,   (c)  Explain,  in  terms  of  the  electron  theory,  why 
they  resemble  one  another  chemically. 

20.  (a)  Draw  a  diagram  of  the  structure  of  a  molecule  of 
potassium  bromide,  KBr.    (b)  Using  your  knowledge  of  the 
electron  theory,  explain  why  the  two  atoms  unite. 

21.  What  is  the  difference  between  electrovalence  and  co- 


22.  How  does  the  electron  theory  explain  why  two  atoms  of 
hydrogen  unite  with  one  atom  of  oxygen  to  form  one  molecule 
of  water? 

23.  How  does  the  electron  theory  explain  the  fact  that  cer- 
tain elements  are  more  active  chemically  than  others? 

24.  In  terms  of  the  electron  theory,  explain  why  some  com- 
pounds are  stable  and  others  unstable. 

25.  Explain  oxidation  and  reduction  in  terms  of  the  electron 
theory.  Use  the  equation  for  the  reducing  action  of  H2  on  CuO 
to  make  your  answer  clearer. 

Group  B 

26.  Study  the  theories  Benjamin  Franklin  held  regarding 
the  nature  of  electricity.  Compare  his  views  with  the  modern 

27.  (a)  What  is  a  cloud  chamber?  (b)  For  what  purpose  is 
it  used?  (c)  What  are  alpha  particles?  (d)  fog  tracks? 

28.  What  evidence  is  there  that  protons  are  always  found 
inside  the  nucleus  of  an  atom? 

29.  Using  the  valence-change  method,  balance  the  follow- 
ing equations: 

a)  HC1  +  Mn02  -»  MnCl2  +  H2O  +  C12  T 

b)  C12  +  H20  ->  2HC1  +  O  T 

c)  KC103  ->  KC1  +  02  T 

d)  S  +  HNO3  -»  H2SO4  +  NO  | 

30.  Why  does  carbon  have  the  2  valences,  -f-2  and  -f-4? 

31.  Explain  why  the  valence  of  a  free  element  must  be  zero. 


1.  Harkins  predicted  the  discovery  of  the  neutron,  and 
Mendeleyeff  predicted  the  discovery  of  several  elements.  Both 
predictions  were  verified.  Such  incidents  in  the  history  of  sci- 
ence are  not  as  rare  as  in  other  fields  such  as  economics  and 
politics.  Give  reasons  for  this. 

2.  Construct  a  model  of  the  sodium  atom  based  on  the  elec- 
tron theory.  You  may  use  wire  for  the  orbits,  a  small  ball  for 
the  nucleus,  and  copper  rods  for  the  electrons. 


AT    LAST! 

The  United  States  knows  that  peace- 
ful power  from  atomic  energy  is 
no  dream  of  the  future.  That  capa- 
bility, already  proved,  is  here  now  — 
today.  I  propose  an  atomic  energy 
agency  under  the  United  Nations  to 
apply  atomic  energy  to  the  needs  of 
agriculture,  medicine.,  and  other 
peaceful  activities  .  .  .  to  provide 
abundant  electrical  energy  in  the 
power-starved  areas  of  the  world. 
Dwight  D.  Eisenhower,  December  8, 

The  dream  of  harnessing  the  energy  of  the  atom.  The  atomic 
bomb  dropped  on  Japan  on  August  6,  1945,  was  one  of  those  incred- 
ible pinnacles  toward  which  fate  drives  unsuspecting  man.  Thou- 
sands of  scientists  who  had  studied  the  atom  and  discovered  the 
electron,  proton,  and  neutron  had  little  idea  of  the  genii  they  were 
uncorking.  Chiefly,  they  were  men  and  women  trying  to  learn  more 
about  the  nature  of  matter,  each  concerned  with  his  own  particular 

A  few  scientists,  it  is  true,  saw  the  possibility  of  someday  learning 
enough  about  the  atom  to  enable  man  to  release  the  tremendous 
forces  locked  within  it.  If  it  could  be  done,  the  energy  locked  within 
a  single  lump  of  coal  would  be  enough  to  drive  a  huge  ocean  liner 
around  the  world.  As  scientists,  they  hoped  for  an  achievement  which 
they  believed  to  be  even  greater  than  the  discovery  of  electricity. 
They  dreamed  of  opening  a  door  to  an  age  of  limitless  power  and 
thus  lifting  the  standards  of  living  of  all  the  peoples  of  the  world. 

When  the  news  of  triumph  finally  came,  it  surprised  even  the  most 
optimistic  of  scientists.  The  great  marvel,  said  President  Truman, 
".  .  .  is  not  the  size  of  the  enterprise,  its  secrecy  or  its  cost,  but  the 




achievement  of  scientific  brains  in  putting  together  infinitely  com- 
plex pieces  of  knowledge  held  by  many  men  in  different  fields  of  sci- 
ence into  a  workable  plan."  The  story  of  some  of  these  men  and  the 
bits  of  knowledge  they  accumulated  has  been  told  in  earlier  chap- 
ters. More  of  the  story  can  now  be  told. 

An  amazing  scientific  discovery  startles  the  world.  Late  in  1895, 
William  Roentgen  (rimt'ge?n)  was  working  in  a  dark  room  with  a 
Crookes  tube  covered  with  black  paper.  He  noticed  that  while  an 
electric  discharge  was  passing  through  the  tube,  a  small  screen  cov- 
ered with  a  chemical,  barium  platinocyanide,  lying  on  a  table  several 
feet  away,  gave  off  a  strange  glow. 

This  was  curious  and  unusual  behavior.  He  thought  it  must  be 
caused  by  rays,  powerful  enough  to  penetrate  not  only  the  glass  of 
the  Crookes  tube  but  several  feet  of  air  as  well.  He  tested  the  pene- 
.  trating  power  of  these  strange  rays  by  producing  them  in  front  of 
several  objects  of  varying  hardness,  including  a  hand  behind  which 
he  had  placed  a  sensitized  photographic  plate.  To  his  astonishment, 
the  film  when  printed  showed  a  hand  with  the  bones  much  darker 
than  the  surrounding  flesh.  He  had,  so  to  speak,  taken  pictures 
through  an  opaque  solid,  a  truly  remarkable  feat! 

These  rays  in  some  ways  acted  like  light,  but  differed  from  light 
in  being  of  much  shorter  wave  lengths  which  could  pass  through 
even  solid  objects.  He  named  these  rays  x-rays.  X-rays  are  produced 
by  the  bombardment  of  matter  by  a  stream  of  rapidly  moving  elec- 
trons, or  cathode  rays.  In  the  cathode  tube,  as  the  cathode  rays  strike 
the  anode,  they  release  a  small  part  of  the  energy  as  x-rays. 

X-ray  machines   (right)  are  used  in  industry  to  detect 
flaws  inside  metal  castings  like  the  crankshaft  (left). 

ml  Klrrtr, 


Another  accident,  and  the  curtain  rises  on  the  drama  of  radio- 
activity. Soon  after  the  discovery  of  x-rays,  another  accident  occurred 
in  the  laboratory  of  Henri  Becquerel  (bek-reT) .  He  was  testing  the 
effect  of  sunlight  on  various  ores,  among  which  he  had,  fortunately, 
included  an  ore  of  uranium. 

Quite  by  chance  he  placed  a  piece  of  the  ore  containing  the  ele- 
ment uranium  upon  a  fresh  photographic  plate  enclosed  in  a  light- 
tight  envelope  lying  on  a  table  in  his  darkroom.  When  he  examined 
this  plate,  he  found  that  it  had  been  changed  under  the  very  spot 
on  which  the  ore  had  rested.  This  was  not  the  sort  of  accident  to 
reach  the  front  page  of  the  newspapers,  as  the  discovery  of  x-rays  had 
done.  But  its  results  were  tremendously  important. 

A  new  world  of  radioactivity  is  discovered.  Becquerel  could  not 
explain  what  had  happened.  He  repeated  the  experiment  with  other 
ores  containing  uranium.  Pitchblende,  he  found,  emitted  similar 
rays,  which  affected  a  photographic  plate  even  more  than  the  other 
uranium  compounds.  He  suspected  some  unknown  element  to  be 
the  cause,  and  asked  Madame  Marie  Curie  (ku-re') ,  a  Polish  girl 
working  as  a  science  teacher  and  research  worker  in  Paris,  to  under- 
take the  isolation  of  this  unknown  element. 

Madame  Curie  and  her  husband,  Pierre,  set  out  to  track  down  the 
cause  of  this  peculiar  behavior  of  pitchblende.  They  boiled  and 
cooked  a  ton  of  this  ore,  sent  to  them  from  the  pitchblende  mines  of 
Austria;  they  filtered  and  separated  out  impurity  after  impurity. 
Years  of  almost  endless  work  passed,  and,  though  they  labored  under 
extreme  difficulties,  Madame  Curie  wrote  years  later:  "It  was  in  that 

Marie  Curie  in  her  laboratory.  For  her 
isolation  of  radium  and  polonium, 
Madame  Curie  was  awarded  the 
Nobel  prize  for  chemistry  in  191 1. 



miserable  shed  of  a  laboratory  that  we  passed  the  best  and  happiest 
years  of  our  lives."  Finally  in  1898  they  succeeded  in  obtaining  a  few 
crystals  of  a  salt  of  a  new  element,  which  they  named  radium. 

The  new  discovery  was  made  public.  A  strange  element  had  been 
discovered  by  a  woman.  Its  salts  shine  in  the  dark  like  tiny  electric- 
light  bulbs  and  emit  heat  continuously.  This  element  is  a  powerful 
poison  —  even  acting  from  a  distance.  It  causes  severe  burns,  and  has 
brought  premature  death  to  a  number  of  scientists  who  have  han- 
dled it  for  long  periods.  It  makes  the  air  around  it  a  good  conductor 
of  electricity. 

Because  radium  captured  the  imagination  of  the  entire  world,  its 
discovery  was  a  great  stimulus  to  further  research.  This  led  to  the  dis- 
covery that  a  number  of  other  elements  resemble  radium  in  their 
ability  to  break  down  and  emit  several  kinds  of  rays.  This  property, 
called  radioactivity,,  is  possessed  by  thorium,  uranium,  polonium, 
radon,  and  several  other  elements. 

In  1902,  Rutherford  and  Soddy,  another  English  scientist,  ex- 
plained the  disintegration  or  breaking  down  of  radium.  Atoms  of 
radioactive  elements,  they  said,  are  not  stable.  They  explode  spon- 
taneously, giving  off  three  types  of  rays:  alpha  and  beta  particles  and 
gamma  rays.  The  gamma  rays  are  similar  to  x-rays;  the  beta  rays  are 
electrons.  Rutherford  later  showed  that  the  alpha  rays  are  electrified 
particles  consisting  of  nuclei  of  helium  atoms.  The  disintegration  or 
decay  of  radium  is  indicated  in  Table  7. 

What  is  the  half -life  of  an  element?  The  half -life  of  an  element  is 
the  time  required  for  the  radioactivity  of  a  given  amount  of  an  ele- 
ment to  decay  to  half  its  original  value.  For  example,  starting  with 
one  gram  of  radium,  it  takes  1620  years  for  half  of  it  to  change  to 
lead.  At  the  end  of  the  next  1620  years,  half  of  the  0.5  gram  which 
is  left  changes  to  lead,  and  this  process  continues  at  the  same  rate. 

TABLE  7. 




at.  wt.  21 8 

half-life  3  min. 

loses  one  helium 

nucleus  of 

at.  wt,  4  and 


at.  wt.  214 

half -life  f9.5  min.' 

loses  one  helium 

nucleus  of 

at.  wt.  4  and 


at.  wt.  210 
half-life  140  days 
loses  one  helium 

nucleus  of 

at.  wt.  4  and 


at.  wt.  206 

a  stable  element 

which  is  the  end 

product  of  radium 


~-!^  radiation 

,90$  molecule 

metal  cylinder  (-) 

tungsten  wire  (  +  )        9|ass  fube 

to  high  potential 
to  R  and  detector  * 

Fig.  35.  Geiger  counter.  An  entering  electron  pro- 
duces a  discharge,  causing  current  to  flow  in  the 
high  resistance  R  until  the  fall  of  potential  across 
R  reduces  the  potential  and  discharge  stops.  The  tube 
is  then  ready  for  the  arrival  of  a  second  electron, 
(left)  Use  of  the  Geiger  counter  to  test  uranium  ore 
for  the  amount  of  radiation. 

Standard  Oil  Company  (XJ.) 

The  Geiger  counter,  detector  of  nuclear  disintegration  products. 

One  type  of  Geiger  counter  consists  of  an  argon-filled  tube  contain- 
ing a  metal  cylinder  and  a  thin  wire.  Between  these  two  electrodes, 
a  very  high  voltage  is  maintained  and  so  adjusted  that  the  tube  is  on 
the  verge  of  discharge.  When  an  electron  enters  the  tube  and  col- 
lides with  a  gas  molecule,  the  tube  is  discharged  and  a  How  of  cur- 
rent is  produced.  This  current  operates  a  headphone,  or  produces  a 
loud  click  or  flash  of  light.  The  greater  the  radiation,  or  the  closer 
the  tube  is  to  the  nuclear  disintegration,  the  greater  the  effect. 

Alpha  particles  are  also  detected  by  the  Geiger  counter.  Neutrons, 
which  have  no  charge,  are  detected  indirectly  with  the  use  of  a  tube 
which  contains  a  boron  compound.  The  boron  nucleus  absorbs  neu- 
trons and  produces  particles  which  may  be  detected.  Gamma  rays 
are  detected  by  the  secondary  electrons  they  produce. 

Alpha  particles,  which  travel  at  10,000-20,000  mi. /sec.,  have  the 
least  penetrating  ability.  About  five  centimeters  ot  air,  a  sheet  of 
paper,  or  a  thin  sheet  (0.1  mm.)  of  aluminum  will  stop  them.  Beta 
particles  (electrons) ,  liberated  at  speeds  more  than  six  times  that  of 
alpha  particles,  require  several  meters  of  air  and  several  millimeters 
of  aluminum  to  absorb  them.  Gamma  rays  have  still  greater  pene- 

Fig.  36.  Alpha,  beta,  and  gamma  rays 
have  different  penetrating  powers  and 
are  affected  differently  by  a  magnetic  field. 

radium  .  ,      ,   v 

salt.  alpha  (a) 

^J_  particles      I 

^Hl .....*..  JW"ii 

••  beta  (fiTT"^ 

lead  Particles 

box  m°9net 

•(        foil 






trating  power  —  several  centimeters  of  aluminum  are  required  to 
stop  them.  Fast  neutrons  have  the  greatest  penetrating  power.  All 
produce  burns.  Radiation  absorption  is  measured  in  roentgens  (r.) . 

Uses  of  radium.  Radium,  and  radon  gas  sealed  in  tubes,  are  used 
in  treating  skin  diseases  and  cancer. 

Considerable  quantities  of  radium  are  used  in  the  detection  of 
(laws  in  castings,  forgings,  and  welds.  Parts  for  aircraft  and  turbine 
casings  are  among  the  many  kinds  of  equipment  tested  in  this  way. 

The  method  of  testing  is  simple.  From  25  to  1000  milligrams  of 
radium  sulfate  are  placed  in  the  center  of  a  circle  of  articles  to  be 
tested.  X-ray  films  are  placed  on  the  backs  of  the  specimens.  The 
penetrating  gamma  rays  from  the  radium  salt  produce  a  shadow- 
graph on  the  film,  quite  like  the  kind  obtained  with  x-rays.  Defects 
as  small  as  0.25  percent  of  the  thickness  of  the  article  can  be  seen 

The  price  of  radium  is  about  $20,000  a  gram.  Carnotite,  an  ore  of 
both  radium  and  uranium,  is  found  in  Utah  and  Colorado.  The  rich 
deposits  of  pitchblende  discovered  on  the  shore  of  remote  Great 
Bear  Lake  in  northern  Canada  in  1931  compete  on  a  favorable  basis 
with  the  huge  deposits  in  the  Belgian  Congo  which  constitute  our 
most  important  source  of  uranium. 

The  attack  upon  the  nucleus  of  the  atom.  The  disintegration  of 
radioactive  elements  indicates  that  the  building  blocks  of  the  atom's 
nucleus  consist  of  neutrons  and  free  protons,  both  of  which  can  be 
emitted.  Helium  nuclei,  which  are  liberated  also  by  radioactive  ele- 
ments, are  themselves  composed  of  neutrons  and  free  protons. 

In  the  effort  to  learn  more  about  the  nucleus,  scientists  began 
bombarding  atoms  with  various  kinds  of  swiftly  moving  projectiles. 
One  of  the  most  important  machines  for  acceleration  of  particles  is 
the  cyclotron  developed  in  1929  by  Ernest  O.  Lawrence.  Some  other 

Brookhavcn  National  Labti 

The  magnet  in  this  giant  cos- 
matron  at  Brookhaven  has 
an  inside  diameter  of  over 
60 feet  and  weighs  2200  tons. 



accelerators  for  smashing  atoms  are  die  betatron,  which  *  accelerates 
electrons,  and  the  synchroton,  cosmotron,  and  bevatron,  all  of  which 
speed  up  protons  to  even  greater  speeds  than  the  cyclotron. 

During  these  researches,  several  discoveries  were  made  —  the  posi- 
tron, artificial  radioactivity,  and  mesons.  The  positron  is  an  ex- 
tremely short-lived  particle,  having  the  same  mass  as  the  electron 
but  opposite  charge.  It  was  discovered  and  named  by  Carl  D.  An- 
derson, of  the  California  Institute  of  Technology,  in  1932.  When  a 
positron  reacts  with  an  electron,  both  particles  disappear  and  gamma 
radiation  is  formed. 

In  the  nucleus  of  the  atom  a  tremendous  energy  lies  like  a  coiled 
spring  caused,  in  part,  by  the  many  protons  which  repel  each  other 
because  they  are  all  of  the  same  electric  charge.  Mesons  may  be  the 
binding  energies  that  hold  the  protons  in  the  nucleus  together. 
There  are  several  types  of  mesons.  The  heavier  meson  (pi-meson)  is 
produced  when  a  nucleus  is  broken  up  during  a  collision.  The 
lighter  meson  (raw-meson)  is  a  decay  product  of  the  heavy  meson. 
Mesons  are  very  short-lived  —  less  than  a  millionth  of  a  second. 

Summary:  Units  of  matter  thus  far  identified.  Let  us  pause  tor  a 
few  minutes  to  list  these  units  of  matter,  their  mass,  charge,  and 
other  related  items  of:  information. 

'«*>''.;  **      *    ti^M'^  <a 

*-.••  ./;.|i 

"  ^&«*"" 


V»***»  - 
t    NAME 

'j^wr""*"*    <•-*        >*~ 

Hfc*:*   *  • 





'•  ^    .*'•>•  t       1 


»    Electron 
J  (6) 

1/1  837  that  of  the 
hydrogen  atom 



In  Crookes  tube,  as 
cathode  rays 

Joseph  John 
Thomson,  1  897 

&  Proton 


Approximately  that 
of  hydrogen 



Stripping  an  electron 
from  hydrogen 

Ernest  Rutherford, 

2  Neutron 
1  (°nl> 

Approximately  that 
of  hydrogen 



Bombarding  beryllium 
with  helium 

James  Chadwick, 

1  Positron 
I  Ut«°) 

Same  as  the 



Radioactive  nitrogen 

Carl  D.  Anderson, 

1  Meson* 

E8.     | 

?  I 

Heavy,  or  pi,  meson 
is  about  285  times 
that  of  electron. 
Light,  or  mu,  meson 
is  about  21  5  times 
that  of  electron. 


nucleus;  as 
cosmic  rays 
in  upper 

Bombarding  atoms 
with  primary  cosmic 
rays  (protons)  or  with 
helium  nuclei  of  300 
million  or  more 
electron-volts  in 

Carl  D.  Anderson 
S.  Neddermeyer 
discovered  the 
mu-meson  in  1  937 
andC.  F.  Powell 
the  pi-meson  in 

5              'This  particle  was  predicted  by  H.  Yukawa,  a  Japanese,  in  1934.  It  was  named  meson, 
s              meaning  Intermediate  porfic/e   It  changes  into  something  else  in  less  than  a  millionth 
$               of  ?  second,  and  it  is  supposed  to  travel  at  a  speed  nearly  that  of  light. 


Ancient  and  modern  alchemy.  As  you  know,  modern  chemistry 
sprang  from  alchemy,  which  was  practiced  for  more  than  20  centu- 
ries. The  chief  goals  of  the  alchemists  were  to  change  the  base  metals, 
such  as  lead  and  iron,  into  gold  and  to  find  an  elixir  that  would  cure 
all  disease.  Although  among  the  alchemists  there  were  many  honest 
enthusiasts,  the  annals  of  their  queer  practice  are  filled  with  accounts 
of  liars  and  charlatans.  In  many  museums  of  Europe  we  can  still  see 
shiny,  yellow  metals  reputed  to  be  gold,  made  by  the  deceptions  of 
the  *  'gold-cooks"  of  European  courts. 

Today  the  alchemists'  dream  of  changing  one  element  into  an- 
other, called  transmutation,  has  come  true.  Radium  changes,  of  its 
own  accord,  into  helium,  lead,  and  other  elements.  Besides  this  natu- 
ral transmutation,  however,  chemists  have  succeeded  in  artificially 
transmuting  many  nonradioactive  elements.  The  first  such  transmu- 
tation was  achieved  in  1919  by  Rutherford,  who  bombarded  nitro- 
gen with  helium  nuclei.  The  nucleus  of  the  nitrogen  atom  was 
changed,  and  one  proton  was  liberated,  the  remaining  nucleus  be- 
coming the  heavy  isotope  of  oxygen  with  atomic  weight  17. 

How  nuclear  reactions  are  written.  Transmutation  or  nuclear  re- 
actions are  like  chemical  equations  and  must  balance.  The  sum  of 
the  subscripts  must  be  the  same  on  both  sides,  and  so  must  the  sum 
of  the  superscripts. 

Rutherford's  transmutation  may  be  expressed  as  follows. 

The  subscripts  (7  +  2-»8+l)  represent  atomic  numbers  and  the 
superscripts  (14  +  4— »  17  +  1)  atomic  weights. 

Recent  experiments  indicate  that  other  elements  may  be  built 
up  from  lighter  elements,  for  example,  carbon  from  beryllium  (a  rare 
metal  lighter  than  aluminum) .  This  transmutation,  (Fig.  37) ,  may 
be  used  to  produce  high-speed  neutrons  by  mixing  powdered  beryl- 
lium with  a  trace  of  a  radium  salt  which  emits  helium  particles. 

Transmutation,  as  you  see,  involves  changes  in  the  nuclei  of  atoms 
rather  than  in  the  shifting  or  sharing  of  electrons  which  produces 
only  chemical  changes. 

Fig.  37.  Transmutation  of  beryllium  into  carbon. 

Beryllium  +      Helium      ^        Carbon 

(at.  wt.  (at.  wt. 



The  search  for  the  key  to  nuclear  energy.  Albert  Einstein,  in  1905, 
advanced  the  idea  that  matter  and  energy  were  really  different  forms 
of  the  same  thing,  and  that  matter  could  be  changed  into  energy  — 
at  least  theoretically.  He  developed  a  mathematical  equation  to  ex- 
press the  conversion  of  matter  into  energy: 

E  =  mc2 

where  E  is  energy  expressed  in  ergs,*  ra  is  mass  expressed  in  grams, 
and  c  is  the  speed  of  light,  expressed  in  centimeters  per  second. 

According  to  this  matter-energy  conversion  equation,  one  pound 
of  matter  (for  example,  one  pound  of  coal  or  of  uranium)  is  equiva- 
lent to  about  11  billion  kilowatthours,  if  completely  changed  into 
energy.  This  is  about  two  and  one-half  times  the  amount  of  electric 
energy  produced  in  an  entire  year  by  the  largest  steam-electric  gen- 
erating plant  in  the  country.  In  burning  the  same  amount  of  coal, 
approximately  four  kilowatthours  of  energy  are  obtained.  In  terms 
of  energy  produced,  oxidation  is,  therefore,  an  extremely  inefficient 

These  ideas,  of  course,  were  all  theory.  However,  a  bit  of  confirma- 
tion came  in  1932.  In  that  year,  Cockcroft  and  Walton,  working  in 
Rutherford's  laboratory,  bombarded  lithium  with  high-speed  pro- 
tons produced  by  accelerating  hydrogen  nuclei  by  Yneans  of  high 
voltages.  They  obtained  helium  (alpha  particles)  with  energies  al- 
most 100  times  as  great  as  the  energy  that  was  used  to  break  the 
lithium  atom.  This  extra  energy  comes  from  the  conversion  of  mat- 
ter into  energy,  in  accordance  with  the  equation  formulated  by  Ein- 
stein, thus: 

Lithium         +  hydrogen    — >  2  helium  +  energy 

3Li7  +  iH1  — »  22He4         +  600,000  electron-volts 

Mass  7.0180  +  mass  1.0076  ->  mass  2(4.0029)  f 

8.0256         ->  8.0058 

Approximate  loss  of  mass  =  0.02 

*  1  erg  ==  1/980  gm.-cm.  of  work  =  approximately  the  energy  required  to  lift 
a  postage  stamp  to  a  height  equal  to  its  thickness.  The  speed  of  light  is 
186,000  mi./sec.  or  30,000,000,000  cm./sec. 

f  Note  that  the  atomic  weight  of  the  isotope  of  the  lithium  used  differs  from 
the  atomic  weight  given  in  the  table  on  page  162,  which  is  an  average  of  the 
atomic  weights  of  all  the  isotopes  of  lithium.  The  hydrogen  here  refers  to  the 
proton,  which  is  slightly  lighter  than  the  hydrogen  atom,  whose  weight  is  given 
in  the  table  on  page  162.  The  weights  given  in  the  table  on  page  162  have  been 
rounded  off  to  three  decimal  places;  hence,  helium  is  shown  there  as  4.003  instead 
of  4.0029  as  in  this  equation. 


However,  the  method  used  by  these  experimenters  was  not  efficient, 
and  there  was  no  great  excitement  over  their  news. 

The  key  is  found.  In  the  meantime,  other  scientists  were  working 
in  this  same  field.  In  1934,  a  young  Italian  physicist,  Enrico  Fermi 
(far'mi) ,  who  later  left  fascist  Italy  to  become  professor  of  physics 
at  Columbia  University,  bombarded  uranium  with  neutrons  and 
thought  he  had  created  a  new  element  No.  93.  Then  four  years  later 
Otto  Hahn  and  F.  Strassman  repeated  Fermi's  experiment  in  Berlin. 
They  bombarded  uranium  with  slow  neutrons  and,  instead  of  pro- 
ducing a  new  artificial  element,  they  obtained  two  other  natural  ele- 
ments and  a  great  deal  of  energy.  Unable  to  explain  what  had  hap- 
pened, Hahn  and  Strassman  nevertheless  published  their  findings. 

Lise  Meitner  (mlt'ner) ,  an  eminent  woman  scientist  working  with 
Hahn,  interpreted  the  results  and  passed  the  information  on  to  Niels 
Bohr,  Nobel  prize  winner  in  physics,  in  Copenhagen.  Dr.  Meitner 
was  forced  to  flee  Germany  by  the  Na/is.  Dr.  Meitner  believed  that 
when  uranium  is  bombarded  by  slow  neutrons,  the  atom  of  uranium 
actually  splits  by  a  process  called  nuclear  fission,  forming  barium  and 
krypton.  But  what  is  even  more  important,  great  quantities  of  en- 
ergy are  released,  perhaps  as  much  as  11  million  kilowatthours  per 
pound  of  uranium.  And,  this  is  only  a  small  part  of  the  energy  that 
would  be  produced  if  all  the  uranium  were  converted  into  energy. 

The  stage  is  set.  Very  soon  after,  a  most  important  conference 
was  held  in  Washington,  D.C.  Atomic  physicists  from  American  col- 
leges and  famous  scientists  from  foreign  nations  were  present.  Niels 
Bohr  was  there,  and  so  was  Enrico  Fermi.  At  this  meeting,  Bohr  and 
Fermi  discussed  the  ideas  of  Meitner.  Bohr  suggested  that  it  was  the 
U-235  in  the  uranium  that  actually  split.  Fermi  suggested  that,  in 
the  fission  of  uranium  by  neutrons,  other  neutrons  might  be  emitted. 
These  emitted  neutrons  could  attack  other  uranium  atoms.  If  this 
were  true,  the  possibility  of  a  chain,  or  self-propagating,  reaction 
that  would  unlock  the  door  to  nuclear  energy  was  near  at  last. 

Brown   Ilroflu 

Enrico  Fermi  (1901-1954),  winner  of  the  1938 
Nobel  prize  for  physics,  played  an  important 
role  in  our  government's  nuclear  research 
program  both  during  and  after  World  War  II. 


Before  the  meeting  in  Washington  was  over,  experiments  to  con- 
firm nuclear  fission  had  begun,  and  confirmation  of  the  emission  of 
neutrons  was  soon  obtained.  By  midsummer  of  1940,  the  important 
facts  regarding  nuclear  fission  had  been  discovered  and  were  known 
by  many  scientists.  And  although  a  chain  reaction  had  not  been  ob- 
tained, its  possibility  was  clear  and  several  methods  of  producing 
it  had  been  suggested.  Then  suddenly,  World  War  II  clamped  tight 
the  door  of  censorship  on  all  research  relating  to  the  release  of 
nuclear  energy.  For  five  years,  the  outside  world  was  kept  in  the  dark. 

Nuclear  energy  unleashed!  With  the  sudden  dropping  of  the  first 
atomic  bomb  on  Hiroshima  in  August,  1945,  the  veil  was  partly 
lifted  on  research  on  nuclear  fission  and  the  production  of  chain 
reactions.  Early  in  1940,  Franklin  D.  Roosevelt  and  Winston 
Churchill  had  pooled  the  efforts  of  British  and  American  scientists 
on  a  research  program,  the  like  of  which  the  world  had  never  seen. 
Its  goal  was  the  release  of  nuclear  energy  for  the  production  of  a 
weapon  with  which  to  win  the  war  against  the  Axis  nations  more 
quickly.  Knowledge  that  research  on  such  a  weapon  was  being  car- 
ried on  in  Nazi  laboratories  compelled  quick,  cooperative  action. 
The  race  was  on  —  the  prize,  the  world.  The  United  States  gov- 
ernment invested  two  billion  dollars  in  ".  .  .  the  greatest  scientific 
gamble  in  history  —  and  won." 

The  term  atomic  energy  has  long  been  used  to  describe  the  tre- 
mendous power  which  is  released  when  nuclear  fission  occurs.  How- 
ever, the  term  nuclear  energy  is  now  preferred  since  it  is  more  truly 
descriptive  of  the  processes  involved. 

A  chain  reaction  from  U— 235.  The  first  controlled  chain  reaction 
was  achieved  on  December  2,  1942,  at  the  University  of  Chicago.  The 
fissionable  material  used  was  pure  U-235  obtained  from  natural 
uranium  ores  containing  a  mixture  of  three  isotopes:  U-234,  U-235, 
and  U-238.  Even  though  only  about  one  part  in  140  of  this  mixture 
is  pure  U-235,  this  isotope  is  used  because  it  is  most  susceptible  to 
nuclear, fission  by  slow  neutrons.  What  happens  in  the  nuclear  fission 
of  U-235  may  be  represented  as: 

U-235  +  neutron  — - »  Ba  +  Kr  +  2  or  3  neutrons  +  energy 

At  Oak  Ridge,  Tennessee,  U-235  was  laboriously  separated  from  its 
other  isotopes  by  an  electromagnetic  method.  A  compound  of  ura- 
nium, UF6,  was  passed  in  the  form  of  a  gas  between  the  poles  of  a 
magnet.  The  lighter  isotope,  U-235,  was  deflected  more  than  its 
heavier  partners  and  thus  separated. 

liberated  neutrons      energy 


56  + 
82  n 

Fig.  38.  A  possible  chain  reaction. 

A  newly  created  element,  plutonium,  for  the  A-bomb.  Few  details 
about  the  A-bombs  exploded  over  Hiroshima  and  Nagasaki  have 
been  released.  It  is  definitely  known,  however,  that  two  fissionable 
elements  were  produced  for  use  in  bombs;  namely,  U-235  and  plu- 
tonium. Plutonium,  which  has  fission  properties  similar  to  U-235,  is 
a  newly  created  element  of  atomic  number  94.  It  was  named  for  the 
planet  Pluto,  which  lies  beyond  Uranus  in  the  solar  system.  Together 
with  another  new  element,  neptunium  of  atomic  number  93,  plu- 
tonium was  first  prepared  in  1940  with  the  aid  of  Lawrence's  cyclo- 
tron by  E.  M.  McMillan  and  P.  H.  Abelson.*  These  were  momentous 

The  formation  of  neptunium  and  plutonium  may  be  represented 
by  the  following  three  equations: 

1)     Uranium  238+ neutron 

uranium  239 

23  mirv 

2)     Uranium  239 ^-neptunium  239  +  electron 

*  Traces  of  these  two  elements  were  later  found  in  some  uranium  ores. 

In  1945,  elements  95,  americium  (Am) ,  and  96,  curium  (Cm) ,  were  obtained 
by  bombarding  plutonium  and  uranium  with  swiftly  moving  helium  nuclei. 

Then  came  elements  97,  berkelium  (Bk) ,  and  98,  californium  (Cf) ,  in  1949 
and  1950.  These  elements  were  named  after  Berkeley  and  California,  the  city  and 
state  in  which  they  were  first  produced.  Elements  99  and  100  were  created  in  1954. 



This  change  occurs  by  the  breaking  down  of  one  neutron  in  the  nu- 
cleus of  U-239  into  one  proton  and  one  electron,  which  escapes. 

2.3  days 

3)     Neptunium  239 *>  plutonium  239  +  electron 

This  change  occurs  by  the  breaking  down  of  one  neutron  in  the 
neptunium  239  nucleus  into  one  proton  and  one  electron,  which 

Plutonium,  in  turn,  becomes  U-235  by  natural  radioactive  disinte- 
gration, or  more  rapidly  by  the  action  of  slow  neutrons  to  which  it 
is  extremely  sensitive.  The  change  is  indicated  by  the  following  nu- 
clear equation: 

Plutonium  —  »  U-235  -f  helium 

What  is  meant  by  "critical  size"?  An  A-bomb  is  set  off  by  suddenly 
bringing  together  two  separate  blocks  of  fissionable  material,  each 
of  which  is  smaller  than  the  critical  size,  but  which  together  form  a 
mass  just  above  this  critical  size.  For  a  bomb  explosion  to  occur,  the 
number  of  neutrons  captured  with  fission  must  be  greater  than  the 
number  of  neutrons  which  escape.  The  number  of  neutrons  which 
escape  depends  on  the  surface  area,  whereas  the  number  captured 
depends  upon  the  volume.  As  the  quantity  of  fissionable  material 
increases,  the  volume  increases  faster  than  the  surface  area.  Critical 
size  means  the  size  at  which  the  neutrons  captured  exceed  the  num- 
ber which  escape  and  fission  occurs. 

Plutonium  is  produced  in  nuclear  reactors  or  "piles."  The  nuclear 
reactor  built  at  Oak  Ridge,  Tennessee,  is  essentially  a  large  cube  of 
graphite  bricks  containing  a  number  of  horizontal  channels  into 
which  is  placed  pure  uranium  in  the  form  of  solid  cylinders  or 
slugs  enclosed  in  aluminum  casings.  Graphite  is  used  to  slow  down 
neutrons  and  is  called  a  moderator.  Heavy  water  is  another  good  mod- 
erator. Slow  neutrons  are  more  effective  in  producing  fission  than 
are  neutrons  that  travel  at  normal  speeds.  The  bricks  are  built  up 
in  layers,  and  since  the  structure  was  built  by  piling  one  layer  of 
bricks  upon  another,  it  is  called  anatomic  pile. 

A  chain  reaction  is  started  with  neutrons  liberated  from  a  bit  of 
beryllium  mixed  with  radium  in  the  center  of  the  pile.  The  concen- 
tration of  neutrons  is  controlled  by  cadmium  or  boron-steel  rods, 
which  absorb  neutrons  easily.  Several  nuclear  reactions  take  place  as 

concrete  shield 
boron  steel  control  rod 

technician  removing 
tubes  containing 

protective  lead  shield 

long  graphite  stringer 

holes  for  aluminum  tubes 

Fig.  39.  Simplified  drawing  of  a 
graphite-moderated  atomic  pile. 

graphite  moderator 




Adapted  from  a  drawing  of  the  Atomic  Enerf/y  Commission 

shown  on  page  187.  When  the  uranium  slugs  are  ready  for  processing, 
they  are  pushed  out  at  the  back  of  the  reactor  and  new  ones  are  fed  in 
at  the  front.  The  slugs  fall  into  tanks  of  water  where  the  U-239  grad- 
ually changes  into  Np-239  and  finally  into  plutonium.  The  slugs  are 
then  dissolved  in  acid,  and  the  plutonium  is  separated  chemically 
from  the  rest  of  the  elements  present.  This  chemical  process  of  sepa- 
ration is  much  easier  than  physical  separation  of  U-235  from  U-238. 

Since  dangerous  radiations  and  radioactive  material  are  produced 
during  these  changes,  all  operations  are  performed  by  remote  control. 
The  whole  pile  is  surrounded  with  several  feet  of  concrete  to  shield 
and  protect  the  operating  personnel.  Several  reactors  are  in  opera- 
tion in  this  country.  The  one  at  Oak  Ridge  is  air-cooled,  while  the 
pile  at  Hanford,  Washington,  is  water-cooled.  Intensive  research  on 
other  coolants  is  now  being  carried  on.  Metals,  such  as  sodium,  in  a 
liquid  state  have  been  found  useful  for  this  purpose. 

The  atomic  pile  liberates  tremendous  amounts  of  heat.  Efforts  to 
build  industrial  nuclear  reactors  which  will  utilize  this  huge  source 


Co  60 

Fig.  40.  Use  of  radioactive  cobalt  for 
the  detection  of  flaws  in  castings. 

defect  in 

Adapted  from  a  drawing  of  the 
Atomic  Energy  Commission 


developed  film 
shows  defect 




of  power  are  already  well  under  way.  Atomic  power  plants  "burning" 
nuclear  fuel  will  supply  electricity,  not  only  in  our  own  country  but 
later  on  also,  it  is  to  be  hoped,  in  those  areas  of  the  world  which  are 
poor  in  coal,  petroleum,  and  natural  water  power.  Nuclear  furnaces 
may  be  built  by  private  industry  with  uranium  supplied  by  the 
Atomic  Energy  Commission.  The  plutonium  manufactured  during 
the  process  will  be  turned  back  to  the  United  States  Government. 

Radioisotopes,  first  produced  artificially  by  Madame  Curie's 
daughter.  In  1934,  Irene  Joliot-Curie  (zho'lyf/)  and  her  husband, 
Frederic  Joliot-Curie,  by  bombarding  boron  with  alpha  particles, 
produced  a  neutron  and  a  radioactive  isotope  of  nitrogen.  Here  was 
another  case  of  modern  alchemy. 

Boron   +  helium  - 
mass  10  +  mass  4  - 

>  radioactive  nitrogen  -f  neutron 
»  mass  13  +  mass  1 

Since  then,  scientists  have  made  more  than  700  new  and  different 
radioisotopes  in  cyclotrons  and  nuclear  reactors.  For  example,  when 
a  bit  of  ordinary  iodine  (atomic  weight  127)  is  placed  in  an  atomic 
pile  where  it  is  bombarded  with  neutrons,  it  changes  to  a  radio- 
isotope  of  iodine  of  mass  131.  The  Atomic  Energy  Commission  sup- 
plies hundreds  of  radioisotopes  to  research  groups  all  over  the  world. 

A  new  "tracer  technique"  uses  radioisotopes.  Research  in  medi- 
cine, biology,  agriculture,  and  many  other  fields  has  been  helped  tre- 
mendously by  this  new  method.  Radioactive  iodine,  for  example,  is 
being  used  in  thyroid  diagnosis  and  therapy.  A  person  suffering  from 
hyperthyroidism  is  fed  with  a  trace  of  sodium  iodide  containing  I131. 
With  the  aid  of  a  Geiger  counter,  the  rate  at  which  this  iodine  com- 
pound collects  in  the  thyroid  gland  can  be  accurately  determined. 

Radioactive  cobalt,  Co"0,  loses  radioactivity  in  five  days  and  is  used 

Atomic  Ei 

Materials  may  be  made 
radioactive  by  exposure 
within  a  nuclear  reactor 
such  as  this  water-boiler 
reactor.  Such  radioisotopes 
have  many  uses  in  indus- 
try, agriculture,  and  medi- 



in  cancer  therapy  as  a  substitute  for  radium  and  x-ray  treatments. 
Radioactive  carbon,  C1*,  is  a  wonderful  tool  in  the  study  of  photo- 
synthesis and  such  problems  of  human  health-  as  sugar  metabolism. 
Radioactive  phosphorus,  P32,  is  used  in  agricultural  research  dealing 
with  the  accumulation,  utilization,  and  action  of  phosphate  fertiliz- 
ers. Industry  is  using  radioisotopes  in  the  improvement  of  steel,  in 
studying  the  action  of  catalysts,  in  measuring  the  flow  of  under- 
ground water,  oil,  and  gas,  and  in  the  detection  of  leaks. 

Nuclear  energy  in  the  future?  So  far,  nuclear  energy  has  been 
used  mainly  as  a  military  weapon,  and  for  research.  No  one  knows 
what  the  peaceful  use  of  nuclear  energy  will  bring.  It  seems  likely 
that  nuclear  science  will  give  higher  standards  of  living  to  all  peoples. 

The  unlocking  of  almost  unimaginable  stores  of  energy  should 
teach  man  important  lessons.  Nuclear  energy  may  transform  the 
world  by  improving  the  health  arid  raising  the  standards  ot  living  of 
millions  of  persons.  But  this  same  instrument  in  the  form  of  a  Hy- 
drogen-bomb can  destroy  civilization  as  the  A-bomb  wiped  out  much 
of  Hiroshima  and  Nagasaki.  Therefore,  the  nations  of  the  world 
must  find  a  way  of  preventing  a  war  with  nuclear  weapons  from  ever 
taking  place. 

The  H-bomb,  based  on  the  fusion  of  the  heavy  isotopes  of  hydro- 
gen into  helium  and  triggered  by  the  100,000,000°C.  temperature  of 
an  A-bomb  explosion,  could  be  made  of  unlimited  size.  This  thermo- 
nuclear reaction  may  be  expressed  as  follows: 

Deuterium  +  tritium  —  >  helium  +  neutron  -f-  energy 
H2         +       H3     —  >    He       + 


The  sudden  conquest  of  nuclear  energy  demonstrated  that  science 
in  a  democracy  is  strong  and  tremendously  creative.  By  constant 
vigilance,  we  must  strive  to  keep  it  so. 

Fig.  41.  Simplified  drawing  of  the  use  of  nuclear  energy  for  generating  electricity. 

Adapted  from  a  drawing  of  the  Atomic  Energy  Comnnsifion 

Reactor  control  console 

Electric  power — « 


A.  >*«,,.  .,.  Turbine 

Reactor  core 
Uranium  rods 







Curie,  Eve.  Madame  Curie.  Garden  City  Publishing  Co., 
New  York,  1943.  In  this  intimate  portrait  of  her  mother,  Eve 
Curie  tells  an  entrancing  story  of  the  discovery  and  isolation 
of  radium. 

Glasstone,  Samuel.  Sourcebook  on  Atomic  Energy.  D.  Van 
Nostrand  Company,  New  York,  1950.  This  book,  written  at 
the  request  of  the  United  States  Atomic  Energy  Commission, 
brings  together  in  readable  narrative  the  important  facts  about 
the  past  history,  recent  status,  and  possible  future  of  nuclear 

Fermi,  Laura.  Atoms  in  the  Family.  The  University  of  Chi- 
cago Press,  Chicago,  1954.  A  simple,  intimate  account  of  the 
events  surrounding  the  conquest  of  a  controlled  nuclear  chain 
reaction,  written  by  the  wife  of  Enrico  Fermi. 

Dean,  Gordon.  Report  on  the  Atom.  Alfred  A.  Knopf,  New 
York,  1953.  An  account  of  the  role  of  nuclear  energy  by  the 
former  head  of  the  U.  S.  Atomic  Energy  Commission. 

Smyth,  Henry  D.  Atomic  Energy  for  Military  Purposes. 
Princeton  University  Press,  1945.  This  is  the  so-called  Smyth 
Report  released  by  the  Army  a  few  days  after  the  first  atomic 
bomb  was  dropped  on  Japan.  It  is  semitechnical  in  nature  and 
not  easy  reading. 


1.  X-rays  consist  of  vibrations  shorter  than  those  of  light. 
They  are  produced  by  bombarding  metals  with  a  stream  of 
rapidly  moving  electrons.  These  x-rays  have  great  penetrating 

2.  Certain  elements  break  down  naturally,  or  disintegrate, 
forming  lighter  elements  including  helium  nuclei.  This  prop- 
erty is  called  radioactivity. 

3.  Radium  gives  off  alpha   particles,   beta   particles,   and 
gamma  rays,  which  are  similar  to  x-rays. 

4.  The  simplest  units  of  matter  thus  far  identified  are  the 
electron,  proton,  neutron,  positron,  and  meson. 

5.  The  positron  has  the  same  mass  as  the  electron;  it  has 
a  positive  charge;  and  it  can  be  formed  during  the  disintegra- 
tion of  radioactive  nitrogen. 

6.  Mesons  have  masses  between  the  mass  of  an  electron  and 
the  mass  of  a  proton;  they  have  either  a  negative  charge,  a  posi- 
tive charge,  or  are  neutral;  they  are  produced  by  the  bombard- 
ment of  atoms  with  cosmic  rays  or  with  rays  of  100-million 
electron-volts  from  the  betatron. 


7.  Modern  alchemy,  or  artificial  transmutation,  is  an  ac- 
complished fact. 

8.  Artificially  radioactive  elements  were  first  obtained  by 
the  Joliot-Curies  by  bombarding  boron  with  alpha  particles. 
Hundreds  of  new  radioisotopes  have  since  been  produced. 

9.  The  age  of  nuclear  energy  was  ushered  in  by  the  con- 
trolled fission  of  U-235  and  plutonium  in  1945. 


Group  A 

1 .  (a)  Who  discovered  x-rays?  (b)  How  are  they  produced? 
(c)  How  do  they  differ  from  light? 

2.  How  did  Becquerel's  discovery  lead  to  the  discovery  of 

3.  List  some  of  the  properties  and  uses  of  Ra. 

4.  What  is  meant  by  the  half-life  of  an  element? 

5.  What  are  the   alpha,  beta,   and  gamma  rays  emitted 
during  the  disintegration  of  Ra? 

6.  Describe  the  construction  and  operation  of  a  Geiger 

7.  Name  five  different  particles  that  have  been  expelled 
from  the  nucleus  of  atoms  during  bombardment. 

8.  Name  several  known  facts  concerning  the  mesons. 

9.  (a)  Has  the  dream  of  transmutation  come  true?  (b)  Ex- 
plain your  answer. 

10.  Show  by  a  diagram  how  Rutherford  changed  N  into  H. 

11.  (a)  Write  the  Einstein  equation  for  the  conversion  of 
mass  and  energy,    (b)  Illustrate  the  meaning  of  this  equation 
in  terms  of  the  change  of  lithium  into  helium. 

12.  What  is  meant  by  a  chain  reaction? 

13.  How  does  artificial  radioactivity  differ  from  the  natural 
radioactivity  of  Ra? 

14.  Describe  some  of  the  events  since  1938  leading  up  to  the 
final  conquest  of  nuclear  energy. 

15.  By  means  of  three  equations,  explain  the  production  of 
plutonium  from  uranium. 

16.  By  means  of  equations,  explain  how  nuclear  energy  was 
released  in  1945. 



17.  Describe  the  construction  of  a  nuclear  pile. 

18.  By  means  of  an  equation,  explain  the  explosion  of  a 
hydrogen  bomb. 

Group  B 

19.  "The  conception  of  the  structure  of  the  atom  makes  it 
possible  for  present-day  scientists  to  explain  the  riddle  of 
transmutation."  Explain  this  statement. 

20.  Make  a  diagram  of  the  heaviest  known  element  showing 
the  composition  of  the  nucleus  and  the  positions  of  all  its 
electrons.  Consult  your  teacher  or  a  recent  edition  of  some  col- 
lege chemistry  textbook. 

21.  Scientists  believe  that  in  releasing  about  four  killowatt- 
hours  of  energy  in  burning  one  pound  of  coal,  a  very  small  part 
of  the  coal  is  converted  into  energy.  Why  has  this  fact  not  been 
proved  by  experiment? 

22.  Describe  briefly  the  future  of  the  peacetime  uses  of 
nuclear  energy. 

23.  (a)  Would  you  buy  stock  in  a  company  organized  to 
exploit  nuclear  energy?  (b)  Give  reasons  for  your  answer. 


1.  Make  a  small  model  of  an  electrical  power  plant  utilizing 
nuclear  fuel.  Consult  your  bibliography  on  page  192. 

2.  Visit  your  dentist  or  doctor  and  ask  him  to  show  you  his 
x-ray  machine.  Make  a  report  on  its  construction  and  opera- 
tion, using  diagrams. 

3.  Consult   your   teacher   of  economics   on   the   question: 
"What  would  the  effect  of  commercial  transmutation  of  iron 
into  gold  be  on  the  financial  structure  of  the  world,  if  it  were 
accomplished  tomorrow?"  Write  a  report  on  this  subject. 

4.  Compare  the  effects  of  the  Industrial  Revolution  and  the 
possible  effects  of  use  of  nuclear  energy.  Ask  your  social  science 
teacher  for  help. 

5.  Write  a  two-page  report  on  the  construction  and  opera- 
tion of  a  cyclotron  or  other  particle  accelerator  now  in  use. 




.  .  .  To  one  man  science  is  a  sacred 
goddess  to  whose  service  he  is  happy 
to  devote  his  life;  to  another  she  is  a 
cow  who  provides  him  with  butter. 
Liebig,  1803-1882 

Gunfire  in  the  American  wilderness  helped  us  to  learn  more 
about  digestion.  In.  1822  at  a  remote  fort  on  Mackinac  Island  be- 
tween Lake  Huron  and  Lake  Michigan,  a  French-Canadian,  Alexis  St. 
Martin,  was  brought  in  tor  medical  treatment.  An  accidentally  dis- 
charged musket  had  sent  a  bullet  through  the  wall  of  his  stomach. 
Dr.  William  Beaumont,  an  army  surgeon,  patched  him  up. 

Despite  great  effort,  it  was  impossible  to  get  the  wound  to  close, 
and  on  healing,  a  flap  covering  an  opening  into  St.  Martin's  stomach 
was  left.  Through  this  opening  Beaumont  could  reach  directly  into 
St.  Martin's  stomach.  Beaumont  got  a  strange  idea.  This  freak  "lid" 
over  the  hole  into  St.  Martin's  stomach  would  enable  him  to  perform 
experiments  to  discover  the  digestive  action  of  the  juices  of  the 
stomach.  St.  Martin  was  agreeable  and  Beaumont  tied  pieces  of  food 
to  a  string,  inserted  them  into  St.  Martin's  stomach  and,  after  several 
hours,  removed  what  was  left  of  the  food. 

In  this  way  Beaumont  gave  science  the  first  accurate  facts  concern- 
ing the  relative  digestibility  of  foods  and  the  composition  of  gastric 
juice.  He  found  gastric  juice  to  contain  a  small  amount  of  hydro- 
chloric acid  (about  0.3  percent) ,  which  helps  to  digest  certain  foods, 
especially  proteins. 




Fig.  42.  Laboratory  prepara- 
tion of  hydrogen  chloride.  Why 
is  the  end  of  the  delivery  tube 
above  rather  than  below  the 
level  of  the  water? 

x  water 

How  hydrogen  chloride  is  prepared  in  the  laboratory.  To  prepare 
hydrogen  chloride  in  the  laboratory,  concentrated  sulfuric  acid  is 
added  to  sodium  chloride  in  a  flask,  as  shown  in  the  illustration 
above.  This  mixture  is  heated,  and  hydrogen  chloride  gas,  which 
is  liberated  readily,  is  collected  by  the  displacement  of  air.  In  this 
chemical  change,  a  double  replacement  occurs,  as  indicated  in  the 
following  equation: 

H2SO4  +  NaCl  -»  HC1  +  NaHSO4  (sodium  bisulfate) 

At  the  outset,  we  must  distinguish  clearly  between  HC1  (hydro- 
gen chloride  gas)  and  HC1  (hydrochloric  acid) .  When  hydrogen 
chloride  gas  is  dissolved  in  water,  hydrochloric  acid  is  obtained. 
Both  are  represented  by  the  same  formula,  but  their  physical  and 
chemical  properties  are  entirely  different. 

How  hydrogen  chloride  is  prepared  commercially.  When  a  jet  of 
hydrogen  is  burned  in  chlorine,  hydrogen  chloride  gas  is  formed  (see 
Chapter  10) . 

H2  +  C12  ->  2HC1 

This  is  one  commercial  method  of  manufacturing  HC1.  A  second 
commercial  method  is  similar  to  the  laboratory  one  but  at  a  some- 
what higher  temperature  producing  Na2SO4  instead  of  NaHSO4. 

2NaCl  +  H2SO4  ->  2HC1  +  Na2SO4 

Chemical  properties  of  hydrochloric  acid.  Hydrochloric  acid  is 
one  of  the  most  common  and  useful  laboratory  chemicals,  or  re- 
agents. Some  of  its  important  properties  are: 

1)  Taste.  Dilute  hydrochloric  acid  has  a  sour  taste. 

2)  Action  on  indicators.  Hydrochloric  acid  reacts  with  a  group  of 
substances  known  as  indicators,  causing  a  color  change.  For  example, 
hydrochloric  acid  turns  blue  litmus  pink.  It  also  turns  reddish-purple 
phenolphthalein  (fe-nol-thal'en)  colorless. 


3)  Action  with  metals.  When  hydrochloric  acid  is  in  contact  with 
most  metals,  a  reaction  takes  place.  Hydrogen  is  liberated  (recall  the 
laboratory  method  for  the  preparation  of  hydrogen) ,  and  chlorides 
are  formed.  Note  that  the  metal  replaces  the  hydrogen  of  the  HC1. 

2HC1  +  Zn  ->  H2  t  +  ZnCl2  (zinc  chloride) 
2HC1  +  Fe  ->  H2  f  +  FeCl2  (ferrous  chloride) 

4)  Action  with  bases,  or  hydroxides  of  metals.  Hydrochloric  acid 
reacts  with  a  base,  forming  a  neutral  compound  that  possesses  the 
qualities  of  neither  acid  nor  base  (bases  are  discussed  in  Chapter  14) . 
Pure  water  is  the  only  other  product  of  this  reaction.  For  example, 
when  hydrochloric  acid  and  sodium  hydroxide  react,  the  products 
are  common  salt  and  water. 

HC1  +  NaOH  -» NaCl  +  H2O 

Why  hydrochloric  acid  is  a  typical  acid.  The  chemical  proper- 
ties of  hydrochloric  acid  are  characteristic  of  the  whole  group  of  com- 
pounds known  as  acids.  We  may  now  define  an  acid  as  a  water  solu- 
tion of  a  compound  with  the  following  characteristic  properties: 

1)  An  acid  has  a  sour  taste.  The  sour  taste,  or  tartness,  of  fruits  is 
caused  by  certain  acids,  such  as  citric  acid,  which  is  found  in  lemons, 
limes,  and  grapefruit. 

2)  An  acid  turns  blue  litmus  pink,  reddish-purple  phenolphthalein 
colorless,  and  acts  with  other  indicators  in  the  same  way  as  hydro- 
chloric acid. 

3)  An  acid  contains  hydrogen  that  can  be  replaced  by  most  metals, 
forming  compounds  known  as  salts.  This  does  not  mean  that  all  com- 
pounds containing  hydrogen  are  acids.  Sugar,  for  example,  contains 
hydrogen,  but  it  is  not  an  acid,  because  its  hydrogen  cannot  be  re- 
placed by  metals.  A  compound  in  which  the  hydrogen  of  an  acid 
has  been  replaced  by  a  metal  is  known  as  the  salt  of  an  acid.  Thus 
sodium  chloride,  NaCl,  is  a  salt  of  hydrochloric  acid,  and  sodium  ni- 
trate, NaNO3,  is  a  salt  of  nitric  acid,  HNO3. 

4)  An  acid  neutralizes  any  base,  forming  water  and  a  salt  whose 
composition  depends  on  both  the  acid  and  the  base  used. 

In  general,  an  acid  consists  of  hydrogen  and  a  nonmetallic  ele- 
ment, or  hydrogen  and  a  radical.  The  hydrogen  of  the  acid  can  be 
replaced  by  a  metal.  Strictly  speaking,  while  certain  substances  show 
acid  properties  only  in  water  solution,  they  are  commonly  called 
acids  even  when  dry.  Thus,  perfectly  dry  H2SO4  is  called  sulfuric 
acid,  even  though  it  exhibits  acid  properties  only  when  it  is  in  a 

Corn  Industries  Research  Foundation,  Inc. 

In  these  converters,  corn  starch  is  changed,  under  heat  and  pressure,  into  corn 
sugar,  or  dextrose.  Hydrochloric  acid  serves  as  a  catalyst  in  the  process.  Because 
the  process  simulates  human  digestion,  dextrose  is,  in  effect,  a  predigested  food 
and  is  readily  assimilated  by  the  body. 

water  solution.  Bases  also  exhibit  their  characteristic  basic  properties 
only  when  water  is  present.  Certain  characteristics  of  acids  or  bases 
in  solution  are  discussed  in  Chapter  16. 

The  general  method  for  the  preparation  of  an  acid.  Sulfuric  acid 
is  one  of  the  most  important  acids,  because  it  is  used  as  a  raw  material 
in  the  manufacture  of  most  of  the  other  acids.  Sulfuric  acid  has  two 
special  characteristics  that  make  it  well  suited  for  this  purpose: 
(1)  its  low  cost,  and  (2)  its  high  boiling  point  (338°C.)  .  The  manu- 
facture of  sulfuric  acid  is  discussed  on  page  309. 

The  preparation  of  hydrochloric  acid  illustrates  the  general 
method  used  in  preparing  acids  from  sulfuric  acid.  First,  a  salt  of 
the  acid  to  be  prepared  is  chosen  as  a  source  of  the  nonmetallic  ele- 
ment of  the  acid.  Common  salt  is  the  least  expensive  and  most  abun- 
dant source  of  chlorine.  Many  other  chlorides  can  be  used,  but  they 
are  more  expensive.  Sulfuric  acid  supplies  the  hydrogen. 

The  salt  of  the  acid  to  be  prepared,  NaCl,  and  sulfuric  acid  are 
heated  together  gently.  Hydrogen  chloride  gas  is  produced  by  the 
reaction  and  is  driven  off  and  dissolved  in  water  in  the  receiving 
vessel,  forming  the  acid.  The  higher  boiling  points  of  the  other  com- 
pounds taking  part  in  the  reaction  prevent  their  vaporizing  and  thus 
keep  them  from  passing  over  into  the  receiving  vessel. 

Many  other  acids  are  manufactured  by  treating  their  least  expen- 
sive and  most  abundant  salts  with  sulfuric  acid.  The  acid  formed  is 




usually  separated  from  the  reacting  substances  by  methods  based  on 
the  differences  in  their  boiling  points. 

Physical  properties  of  hydrochloric  acid.  Hydrochloric  acid  is  a 
colorless  liquid,  heavier  than  water.  That  is,  its  specific  gravity  is 
greater  than  one.  It  possesses  an  irritating  odor.  Both  the  boiling- 
point  and  the  specific  gravity  of  hydrochloric  acid  are  determined  by 
the  weight  of  hydrogen  chloride  gas  dissolved  in  the  water.  Hydro- 
chloric acid  containing  20  percent  hydrogen  chloride  gas  by  weight 
boils  at  110°C.  Impure  hydrochloric  acid  is  called  muriatic  acid  and 
is  usually  yellow  in  color.  In  this  form  it  was  known  for  many  years 
before  Priestley,  in  1772,  first  isolated  pure  hydrochloric  acid. 

Hydrochloric  acid  cleans  metals.  Before  coating  metals,  such  as 
iron  and  steel,  with  plates,  films,  or  coatings  of  other  metals,  includ- 
ing chromium,  silver,  tin,  and  zinc,  the  surface  of  the  metal  must  be 
clean  and  free  of  oxides.  Removing  oxides  and  otherwise  cleaning 
the  surface  of  a  metal  to  be  plated  or  coated  is  a  process  known  com- 
mercially as  pickling.  One  of  the  chief  industrial  uses  of  hydro- 
chloric acid  is  in  the  pickling  of  metals,  especially  before  coating 
with  tin  in  tinning,  or  xinc  in  galvanizing,  or  with  the  materials  that, 
after  firing,  result  in  enamelware. 

Small  quantities  of  hydrochloric  acid,  usually  muriatic  acid,  are 
used  in  removing  rust  stains  from  vitreous  washbasins  and  lavatories. 
Plumbers  often  use  muriatic  acid  as  a  flux  before  soldering. 

Racks  of  sheet  steel  emerge  from  the  pickler.  The  worker  in  the  foreground  is 
dipping  the  sheets  in  water  to  remove  the  acid. 


Hydrochloric  acid  is  used  in  making  other  chemicals.  Chlorides 
of  many  metals,  including  magnesium  chloride,  aluminum  chloride, 
and  zinc  chloride,  are  made  by  the  reaction  of  hydrochloric  acid  and 
a  carbonate  or  oxide  of  the  metal.  For  the  most  part,  such  chemicals 
are  of  very  high  quality,  and  are  used  chiefly  by  manufacturing  chem- 
ists and  drug  houses,  and  by  druggists.  Zinc  chloride  is  used  to  im- 
pregnate wood  to  prevent  decay,  in  soldering,  and  in  flame-proofing. 

Silver  chloride,  AgCl,  one  of  the  several  light-sensitive  silver  com- 
pounds used  to  coat  photographic  film  may  be  made  by  the  reaction 
of  silver  nitrate,  AgNO3,  and  hydrochloric  acid.  In  the  manufacture 
of  glucose  from  starch,  hydrochloric  acid  acts  as  a  catalytic  agent.  It 
is  used  also  in  large  quantities  in  the  manufacture  of  glue  and  gela- 
tin, in  the  purification  of  boneblack,  and  in  the  processing  of  textiles. 

Physical  properties  of  hydrogen  chloride  gas.  Hydrogen  chloride 
gas  is  colorless,  heavier  than  air,  and  has  a  sharp,  penetrating  odor. 
It  is  extremely  soluble  in  water.  If  a  test  tube  of  hydrogen  chloride 
gas  is  placed  mouth  downward  in  water,  the  water  will  rise  almost 
to  the  top,  as  the  gas  dissolves.  Water  dissolves  about  500  times  its 
own  volume  of  this  gas  under  normal  laboratory  conditions.  The 
gas  can  be  liquefied  and  solidified,  just  as  all  other  gases. 

Chemical  properties  of  hydrogen  chloride  gas.  Hydrogen  chloride 
does  not  show  acid  properties  unless  it  is  dissolved  in  water  or  unless 
some  water  vapor  is  present.  When  dry,  it  is  completely  inactive.  Its 
attraction  for  water  is  so  great  that  it  forms  a  cloud  or  mist  in  moist 
air  of  tiny  droplets  of  hydrogen  chloride  solution. 

This  property  may  be  used  as  a  test  for  hydrogen  chloride  gas.  If 
you  blow  across  the  mouth  of  a  bottle  containing  concentrated  hy- 
drochloric acid,  a  mist  will  form.  This  mist  is  caused  by  the  moisture 
in  your  breath  combining  with  the  hydrogen  chloride  vapor  that 
rises  from  the  bottle.  Hydrogen  chloride  does  not  burn. 

Preparation  and  properties  of  hydrofluoric  acid,  HF.  The  prep- 
aration of  hydrofluoric  acid  follows  the  general  method  for  making 
an  acid.  Sulfuric  acid  reacts  with  calcium  fluoride,  CaFo,  the  most 
common  salt  of  hydrofluoric  acid,  forming  hydrofluoric  acid.  This  is 
the  method  used  by  Scheele  when  he  first  prepared  it  in  1771.  Cal- 
cium fluoride  is  the  chief  constituent  of  the  mineral  fluorspar,  found 
in  several  parts  of  the  United  States. 

H2SO4  +  CaF2  ->  2HF  |  +  CaSO4 

Pure  anhydrous  hydrogen  fluoride  is  a  colorless  gas  which  boils  at 
room  temperature  (19.5°C.) .  It  is  deadly  if  inhaled.  It  dissolves 
in  water,  forming  a  colorless  acid  that  vaporizes  at  low  temperatures. 

Calibration  marks  are  carefully 
scratched  through  the  wax  coat- 
ing on  a  graduated  glass  cylinder 
prior  to  exposing  it  to  hydrofluoric 

Corning   (Hass   Works 

Such  an  acid  is  called  a  fuming  acid.  Since  hydrofluoric  acid  re- 
acts with  glass,  quartz,  and  most  metals,  it  is  distilled  in  dishes  made 
of  lead  and  must  he  kept  in  bottles  made  of  polyethylene  or  other 
plastics  with  which  it  does  not  react.  This  acid  produces  powerful 
burns  by  poisoning  the  tissues. 

How  glass  is  etched  with  hydrofluoric  acid.  Etching  is  the  biting 
out  of  particles  of  glass  or  metal  by  means  of  chemicals.  Hydrofluoric 
acid  etches  glass  because  it  unites  with  the  silicon  dioxide,  SiO2,  of 
the  glass,  forming  silicon  tetrafluoride,  SiF.,,  which  is  a  gas. 

SiO2  +  4HF  ->  SiF4 1  +  2H2O 

In  etching  glass  articles,  such  as  thermometers,  electric-light  bulbs, 
and  windows,  the  surface  is  completely  covered  with  wax  and  the 
design  to  be  etched  is  scratched  through  the  wax.  The  object  is  then 
brought  into  contact  with  the  vapor  of  hydrofluoric  acid.  When  the 
action  on  the  exposed  glass  has  gone  as  far  as  necessary,  the  object  is 
removed  from  contact  with  the  vapor.  In  frosting  the  inside  of  an 
electric-light  bulb,  a  small  amount  of  hydrofluoric  acid  is  poured 
into  the  bulb,  shaken  for  an  instant,  and  poured  out,  and  the  bulb 
is  thoroughly  rinsed.  It  is  also  used  in  making  Freon  refrigerants,  and 
as  a  catalyst  in  the  manufacture  of  high  octane  gasoline. 

The  other  halogen  acids.  Theoretically,  both  hydrobromic  acid. 
HBr,  and  hydriodic  acid,  HI,  may  be  prepared  by  the  general  method 
used  in  preparing  acids. 

H2SO4  +  KBr  -»  HBr  |  +  KHSO4 
H2SO4  +  Nal  ->  HI  j  +  NaHSO4 




When  first  formed  they  are  colorless  gases  with  strong  irritating 
odors,  but  are  almost  immediately  oxidized  by  the  oxygen  of  the  air, 
forming  free  bromine  and  free  iodine.  In  water  solution,  hydro- 
bromic  and  hydriodic  acids  are  unstable;  on  exposure  to  air  they 
decompose  as  a  result  of  oxidation. 

4HBr  +  O2  -»  2H2O  +  2Br2 

Heat  of  formation  and  chemical  stability.  When  hydrogen  and 
chlorine  react,  forming  hydrogen  chloride,  a  great  deal  of  heat  is 
produced.  A  reaction  in  which  heat  is  liberated  is  called  an  exother- 
mic reaction  (ex  —  out,  therme  —  heat) .  Exothermic  reactions  con- 
tinue after  they  are  first  started  by  external  heat.  On  the  other  hand, 
when  hydrogen  and  iodine  unite,  heat  is  continuously  absorbed  and 
heat  must  be  added  if  the  reaction  is  to  continue.  A  reaction  in 
which  heat  is  absorbed  is  an  endothermic  reaction. 

The  number  of  calories  of  heat  absorbed  or  liberated,  during  the 
formation  of  a  mole  (see  page  130)  of  an  element  or  compound, 
is  called  its  heat  of  formation.  A  substance  which  liberates  heat  dur- 
ing its  formation  is  said  to  have  a  positive  heat  of  formation;  a  sub- 
stance which  absorbs  heat  is  said  to  have  a  negative  heat  of  forma- 
tion. A  compound  such  as  sodium  chloride,  NaCl,  whose  positive 
heat  of  formation  is  very  great  (97,800  calories)  iswery  stable.  Hy- 
drobromic  acid,  HBr,  whose  positive  heat  of  formation  is  small  (8400 
calories)  is  somewhat  unstable.  Hydriodic  acid,  HI,  which  has  a 
negative  heat  of  formation,  is  very  unstable. 

A  knowledge  of  the  heats  of  formation  is  very  useful  to  chemists. 
For  example,  we  can  tell  whether  a  certain  compound  will  form  and 
how  easily  it  can  be  prepared.  A  compound  formed  by  replacement 
has  a  higher  positive  heat  of  formation  than  the  compound  from 
which  it  is  formed  and  hence  is  more  stable. 





By  similar  reasoning  we  can  see  why  bromine  will  not  replace  chlo- 
rine from  KC1. 




Clendening,  Logan.  The  Human  Body*  pp.  74-75.  Alfred  A. 
Knopf,  New  York,  revised  edition.  1945.  An  amusing  story  of 
Dr.  William  Beaumont's  experiments  with  Alexis  St.  Martin. 

Jaffe,  Bernard.  Men  of  Science  in  America,  pp.  157-158. 
Simon  and  Schuster,  New  York,  1944.  The  dramatic  story  of 
the  pioneer  investigations  of  the  American  doctor,  William 
Beaumont,  and  of  John  R.  Young,  which  resulted  in  our 
knowledge  of  the  presence  of  hydrochloric  acid  in  the  gastric 

Lowry,  T.  M.  Historical  Introduction  to  Chemistry, 
pp.  12-16.  The  Macmillan  Co.,  London,  1915.  The  discovery 
of  the  common  acids  and  their  chemical  and  physical  proper- 
ties are  reviewed. 


1.  An  acid  is  a  substance  whose  water  solution    (1)  has  a 
sour  taste,   (2)  turns  blue  litmus  pink,   (3)  contains  hydrogen 
that  can  be  replaced  by  each  of  many  metals  with  the  forma- 
tion of  a  salt  and  the  liberation  of  hydrogen,  and    (4)  neu- 
tralizes bases,  forming  a  salt  and  water. 

2.  Chemical  reactions  in  which  heat  is  liberated  are  called 
exothermic  reactions;   those   in  which  heat  is  absorbed  are 
called  endothermic  reactions. 

3.  The  heat  of  formation  of  a  compound  is  the  number  of 
calories  of  heat  liberated  or  absorbed  in  the  formation  of 
1  gram-molecular  weight  of  the  compound. 

4.  Compounds  with  high  positive  heats  of  formation  are 
stable;  compounds  with  low  positive  heats  of  formation  are  un- 
stable; compounds  with  negative  heats  of  formation  are  very 


Group  A 

1.  Describe  the  laboratory  method  of  preparing  HC1. 

2.  (a)  What  are  the  two  methods  of  preparing  hydrogen 
chloride  commercially?  Write  an  equation  for  each  method. 
(b)  What  type  of  chemical  reaction  does  each  method  repre- 


3.  Give  the  properties  of  HC1  as  to  (a)  taste,  (b)  action  on 
indicators,   (c)  action  on  metals,   (d)  action  on  oxides,   (e)  ac- 
tion on  hydroxides  of  metals. 

4.  (a)  What  is  a  salt?  (b)  Give  three  examples. 

5.  (a)  What  are  the  salts  of  HC1  called?   (b)  Name  three. 

6.  (a)  What  are  salts  of  HNO3  called?   (b)  Name  one. 

7.  (a)  What  is  an  acid?  (b)  Name  five  acids. 

8.  Make  a  list  of  the  properties  of  hydrogen  chloride  gas. 

9.  What  is  the  percentage  of  hydrogen  chloride  in  a  solu- 
tion that  has  a  fixed  boiling  point  of  110°C.? 

10.  What  weight  of  pure  salt  is  needed  to  prepare  292  Ib.  of 
hydrochloric  acid  containing  15  percent  HC1  by  weight? 

1  1  .  Compare  the  solubility  of  hydrogen  chloride  in  H2O 
with  the  solubility  of  air,  O2,  H2,  CO2,  and  N2. 

12.  What  causes  the  fuming  of  hydrogen  chloride  in  moist 

13.  What  test  or  tests  would  you  use  in  showing  the  pres- 
ence of  HC1  in  a  liquid? 

14.  (a)  What  are  the  uses  of  HG1?  (b)  What  process  in  the 
human  body  depends  in  part  upon  HC1? 

15.  (a)  State    the    general    method    of    preparing    acids. 
(b)  Why  is  H2SO4  so  generally  used? 

16.  Scheele  was  the  first  to  prepare  HF.  His  method  is  still 
used  today.  Describe  it. 

17.  By  means  of  an  equation,  explain  the  etching  action 
of  HF. 

18.  Explain  how  HF  is  stored. 

19.  Compare  the  properties  of  the  other  halogen  acids  with 
those  of  HC1. 

20.  (a)  What  is  a  calory?   (b)  What  is  heat  of  formation? 

21.  (a)  What  is  the  difference  between  an  exothermic  re- 
action and  an  endothermic  reaction?   (b)  Give  an  example  of 

22.  Explain  the  relationship  between  the  heat  of  formation 
and  the  stability  of  a  substance. 

23.  The  positive  heat  of  formation  of  H2S  is  2730  calories. 
(a)  Is  H.,S  a  stable  compound?  (b)  Explain. 


24.  Write  balanced  equations  for  the  following: 

a)  Copper  oxide  +  hydrochloric  acid  — * 

b)  Hydriodic  acid  4-  oxygen  — > 

c)  Sulfuric  acid  +  sodium  bromide  — » 

Group  B 

25.  What  is  the  effect  of  chlorine  water  on  litmus? 

26.  When  HC1  is  boiled,  what  passes  off?  Explain. 

27.  A  bottle  of  HI  turns  brown.  Explain. 

28.  What  chemical  tests  would  you  use  in  identifying  each 
of  the  halogen  acids? 

29.  Why  is  muriatic  acid  used  in  the  soldering  of  metals? 

30.  Water  may  be  considered  an  acid.  Explain. 

31.  Why  is  the  general  method  for  preparing  acids  not  used 
in  the  preparation  of  HI? 

32.  In  the  electrolysis  of  HC1,  the  volumes  of  C12  and  H2 
collected  are  not  the  same.  Explain. 


1.  Take  home  some  strips  of  litmus  paper  and  test  various 
foods  and  other  substances  found  at  home  for  the  presence  of 
acids.  Report  your  findings  to  the  class.  Make  a  list  of  common 
foods  that  react  to  your  test  as  acids;  as  bases.  What  do  you 
conclude  from  your  tests? 

2.  Prepare  a  report  on  the  experiments  performed  by  Beau- 
mont with  Alexis  St.  Martin  which  dealt  with  the  discovery  of 
HC1  in  gastric  juice. 

3.  Consider  carefully  the  quotation  at  the  beginning  of  this 
chapter.  Prepare  a  report  for  the  class  on  the  meaning  of  this 
quotation.  Illustrate  with  examples  from  the  lives  of  great 
scientists  and  inventors. 




.  .  .  Rouelle  [the  teacher  of  Lavoi- 
sier], to  whom  we  owe  the  term  base, 
described  how  natural  salts  had  been 
restricted  at  first  to  salts  formed  by 
the  union  of  acids  with  alkalies 
which  are  soluble  in  water,  and  im- 
part on  the  tongue  a  saline  taste. 
T/M.  Lowry,  1915 

How  sodium  hydroxide  is  prepared  in  the  laboratory.  The  lab- 
oratory preparation  of  hydrogen  by  the  action  of  sodium  on  water 
was  discussed  in  Chapter  3.  Each  sodium  atom  replaces  one  of  the 
atoms  of  hydrogen  in  the  water  molecule. 

2Na  +  2HOH  ->  SNaOH  +  H2 1 

After  the  water  evaporates,  a  white  solid,  sodium  hydroxide,  is  left. 

Chief  properties  of  sodium  hydroxide.  Sodium  hydroxide  is  a 

common  and  useful  substance.  Some  of  its  important  properties  are: 

1)  Action  on  indicators.  Sodium  hydroxide  turns  pink  litmus  blue 
and  colorless  phenolphthalein  reddish-purple. 

2)  Action  with  acids.  Sodium  hydroxide  reacts  with  an  acid,  form- 
ing a  salt,  thus  causing  the  original  properties  of  the  acid  to  disap- 
pear along  with  its  own.  For  example,  sodium  hydroxide  unites  with 
hydrochloric  acid,  forming  common  salt  and  water.  The  presence  of 
the  salt  may  be  proved  by  evaporating  the  water  and  tasting  the 
solid  left  behind. 

3)  Feel.  A  water  solution  of  sodium  hydroxide  has  a  slippery, 
soapy  feeling. 



Why  sodium  hydroxide  is  a  typical  base.  The  chief  properties  of 
sodium  hydroxide  are  characteristic  of  the  whole  group  of  com- 
pounds known  as  bases.  We  may  now  define  a  base  as  a  compound 
(1)  that  contains  the  hydroxyl  group  (OH) ,  (2)  whose  water  solu- 
tion is  soapy  to  the  touch,  (3)  whose  water  solution  turns  pink 
litmus  blue  and  colorless  phenolphthalein  reddish-purple,  and  (4) 
whose  water  solution  reacts  with  acids,  forming  water  and  a  salt  of 
the  acid  used. 

In  general,  then,  a  base  consists  of  a  metallic  element  or  radical  and 
one  or  more  hydroxyl  groups  (OH)  . 

Strong  bases,  such  as  sodium  hydroxide  and  potassium  hydroxide, 
are  often  known  as  alkalies.  Three  other  important  bases  are  cal- 
cium hydroxide  (slaked  lime) ,  ammonium  hydroxide  (ammonia 
water) ,  and  magnesium  hydroxide  (milk  of  magnesia) . 

How  bases  react  with  acids.  When  a  base  reacts  with  an  acid,  the 
hydrogen  of  the  acid  combines  with  the  hydroxyl  radical  of  the 
base,  forming  water,  and  the  metal  or  metallic  radical  of  the  base 
unites  with  the  nonmetal  or  acid  radical  of  the  acid,  forming  a  salt. 
Such  a  chemical  change  is  called  neutralization.  All  neutralization 
reactions  are  double  replacements. 

HC1  +  NaOH  -»  HOH  +  NaCl  (a  salt) 
HNOa  +  KOH  -*  HOH  +  KNO3  (a  salt) 
H2SO4  +  Ca(OH)2  -»  2HOH  +  CaSO4  (a  salt) 
HBr  +  NH4OH  ->  HOH  +  NH4Br  (a  salt) 

Neutralization  is  the  reaction  of  an  acid  with  a  base,  forming  wa- 
ter and  a  salt.  A  salt,  then,  is  a  compound  made  up  of  a  metal  or  a 
metallic  radical  and  a  nonmetal  or  an  acid  radical. 

Titration  and  the  use  of  molar  and  normal  solutions.  It  is  some- 
times necessary  to  know  how  much  of  an  acid,  a  base,  or  a  salt  is 
present  in  a  solution.  For  example,  we  may  want  to  know  how  much 
acetic  acid  is  present  in  a  given  sample  of  vinegar.  One  way  is  to 
neutralize  a  given  volume  of  this  liquid  with  a  solution  of  a  base 
whose  composition  is  known.  This  process  of  determining  the 
strength  of  an  acid  or  a  base  with  the  help  of  neutralization  reaction 
is  called  titration.  It  is  carried  out  in  long  tubes  (burettes)  which 
are  measured  off  in  milliliters  so  that  volumes  can  be  read  off  di- 
rectly. A  definite  volume  of  the  acid  solution  to  be  tested  is  neutral- 
ized with  a  standard  solution  of  a  base,  that  is,  one  whose  composi- 
tion is  known.  The  point  at  which  neutralization  is  reached  (end 
point)  is  determined  by  the  use  of  an  indicator  such  as  phenolphtha- 
lein which  suddenly  changes  color. 



Fig.  43.  Titration.  The  degree  of  acidity  of 
any  acid  solution  is  determined  by  measur- 
ing the  amount  of  normal  basic  solution 
required  to  neutralize  it.  The  process  is 
reversed  to  find  the  strength  of  any  basic 

-* —  burette 


tT      *%&(* 

„*         -" 

There  are  two  kinds  of  standard  solutions  used,  namely,  molar 
solutions  and  normal  solutions.  One  liter  of  a  molar  solution  con- 
tains one  mole  or  gram-molecular  weight  of  dissolved  substance.  For 
example,  a  molar  solution  of  Nad  contains  23  +  35.5,  or  58.5,  grams 
of  NaCl  in  one  liter  of  this  solution.  A  molar  solution  of  Ca  (OH)  2 
contains  74  g. 

One  liter  of  a  normal  acid  solution  contains  one  gram  of  replace- 
able hydrogen.  One  liter  of  a  normal  basic  solution  contains  17  grams 
of  the  OH  group,  and  one  liter  of  a  normal  salt  solution  contains  the 
equivalent  of  one  gram  of  replaceable  hydrogen.  For  example,  one 
liter  of  a  normal  solution  of  H2SO4  contains  98  -4-  2,  or  49,  grams 
H2SO4  since  this  acid  contains  two  replaceable  hydrogen  atoms.  One 
liter  of  a  normal  Ca  (OH)  2  solution  contains  74  -:-  2  grams  Ca  (OH)  2. 
One  liter  of  a  normal  A1C18  solution  contains  133.5  -4-  3  grams  A1C1S 
since  this  salt  contains  three  hydrogen  equivalents.  Decimal  frac- 
tions are  used  in  referring  to  both  molar  (M)  and  normal  (N)  solu- 
tions. Thus  a  0.1M  solution  of  HC1  contains  36.5  -~  10,  or  3.65, 
grams  HC1  per  liter  of  solution. 


1.  How  much  NaCl  per  liter  of  solution  does  a  0.3M  solu- 
tion of  NaCl  contain? 

2.  How  much  acid,  base,  or  salt  per  liter  of  solution  do  the 
following    solutions    contain:    0.1N    AgNO3,    0.3N    MgBr2, 
1.5N  H3PO4,  0.125N  CuCl2? 

3.  How  would  you  prepare  the  following  solutions: 

0.8N  HNO3,  3.0N  KOH,  0.25N  Mg(OH), 

1.3N  A1(OH)3? 


Heat  of  neutralization.  All  neutralization  reactions  are  exothermic. 
When  any  strong  acid  and  strong  base  react,  a  definite  quantity  of 
heat  is  liberated,  13,700  calories  per  gram-molecular  weight,  or  mole, 
of  water  formed.  This  heat  is  called  heat  of  neutralization.  When 
36.5  grams  of  hydrochloric  acid  react  with  40  grams  of  sodium  hy- 
droxide, 18  grams  of  water  and  58.5  grams  of  sodium  chloride  are 
formed  with  the  liberation  of  13,700  calories.  In  the  case  of  the  ac- 
tion of  sulfuric  acid  on  potassium  hydroxide,  2  X  13,700  calories  are 
produced,  because  two  moles  of  water  are  formed, 

H2SO4  +  2KOH  -»  K2SO4  +  2H2O 

Important  uses  of  neutralization.  After  petroleum  has  been  puri- 
fied with  the  aid  of  sulfuric  acid,  the  excess  acid  is  removed  by  neu- 
tralizing it  with  a  base,  usually  sodium  hydroxide.  Much  of  the 
soil  in  the  United  States  is  sour  (acid) .  Four-fifths  of  the  cultivated 
land  in  the  central  western  states  is  sour  and  therefore  not  fully 
productive.  These  acid  soils  may  be  neutralized  by  the  addition  of 
lime,  CaO,  which  combines  with  water,  forming  a  base.  Calcium  car- 
bonate, CaCOa,  is  used  also.  Soil  with  excess  lime  is  neutralized  with 
ammonium  sulfate,  (NH4)  2SO4. 

The  destructive  effect  of  acid  stains  or  burns  may  be  minimized 
by  the  prompt  application  of  either  a  weak  base  or  sodium  bicar- 
bonate, which  neutralizes  the  effect  of  the  acid.  Similarly,  an  alkali 
burn  is  treated  by  application  of  a  mild  acid,  such  as  boric  acid  or 
vinegar.  An  excessive  acid  mouth  condition  is  sometimes  treated  by 
using  small  amounts  of  milk  of  magnesia,  Mg  (OH)  2,  a  mild  base,  or 

Normal  blood  is  slightly  alkaline.  Alkaline  reserve  refers  to  the 
amount  of  base  present  as  bicarbonate  in  the  blood.  Even  slight 
changes  in  the  normal  alkalinity  of  the  blood  result  in  serious  body 
disturbances.  Such  disturbances  may,  in  some  cases,  be  corrected  by 
using  neutralizing  chemicals  administered  by  a  physician. 

Existence  of  an  alkaline  reserve  (a  scientific  fact)  is,  unfortunately, 
used  by  some  advertisers  as  a  reason  for  selling  to  the  public  huge 
quantities  of  "alkalizers"  to  dose  ailments  that  are  best  treated  by 
other  methods  administered  by  a  competent  physician.  This  is  an 
excellent  example  of  the  abuse  of  scientific  knowledge,  against  which 
intelligent  persons  must  always  be  on  guard. 

Physical  properties  of  sodium  hydroxide.  The  white  solid,  sodium 
hydroxide,  is  very  soluble  in  water.  On  exposure  to  moist  air,  it  ab- 
sorbs large  quantities  of  water  and  changes  into  a  pool  of  sodium 
hydroxide  solution.  This  property  of  deliquescence  makes  sodium 

hydroxide  useful  as  a  drying  agent.  Usually  it  is  sold  in  pellets  that 
must  be  kept  in  well-sealed  bottles  which  exclude  moisture. 

Chemical  properties  of  sodium  hydroxide.  In  addition  to  typical 
properties  of  bases  already  mentioned,  sodium  hydroxide  has  other 
chemical  properties.  When  exposed  to  air,  it  unites  with  carbon  di- 
oxide, forming  sodium  carbonate. 

2NaOH  +  CO2  ->  Na2CO3  +  H2O 

The  common  name  for  sodium  hydroxide  is  caustic  soda,  given  to 
it  because  of  its  caustic,  or  burning,  action  upon  the  skin.  It  dis- 
solves wool  but  has  little  effect  upon  cotton,  rayons,  or  nylons.  Be- 
cause of  this  fact,  it  is  used  in  determining  the  amount  of  wool  in  a 
cotton-wool  mixture. 

Potassium  hydroxide  (caustic  potash}  closely  resembles  sodium  hy- 
droxide in  both  chemical  and  physical  properties. 

Sodium  hydroxide  helps  clothe  us.  Sodium  hydroxide  is  one  of 
the  most  useful  compounds  known.  It  serves  man  in  many  ways, 
chief  among  which  is  in  the  manufacture  of  rayon  and  cellulose 
films  similar  to  cellophane.  Approximately  one-fifth  of  the  3  million 
tons  of  sodium  hydroxide  produced  in  a  recent  normal  year  in  this 
country  was  consumed  in  this  way.  The  chemistry  of  the  production 
of  rayon  and  cellulose  films  is  discussed  in  detail  in  Chapter  37. 

Many  fabrics  are  made  of  cotton  that  has  been  treated  with  so- 
dium hydroxide.  When  cotton  fibers  are  placed  in  a  solution  of 
sodium  hydroxide,  they  lose  part  of  their  natural  twist  and  acquire 
a  gloss  that  is  considered  very  desirable  by  many  persons.  Cotton  so 
treated  is  known  as  mercerized  cotton. 

Sodium  hydroxide  helps  keep  us  clean.  Until  very  recently  sodium 
hydroxide  and  potassium  hydroxide  were  used  in  making  almost  all 
the  cleansing  agents  for  both  industrial  and  home  use.  When  sodium 


hydroxide  or  potassium  hydroxide  reacts  with  a  fat,  soap  and  glycerin 
are  formed.  Each  year  the  soap  industry  uses  thousands  of  tons  of 
sodium  hydroxide  and  considerable  amounts  of  potassium  hydroxide 

Another  important  use  of  sodium  hydroxide  is  in  the  form  of  he 
which  contains  about  94  percent  sodium  hydroxide.  It  is  a  useful 
household  cleansing  agent  because  it  dissolves  grease.  Lye  produces 
pamlul  burns  if  it  comes  in  contact  with  the  skin.  It  should  be  used 
cautiously  and  stored  out  of  the  reach  of  small  children 

When  kitchen  or  bathroom  plumbing  becomes  clogged,  a  strong 
solution  of  lye  is  sometimes  poured  down  the  pipes.  In  this  way 
greasy,  fatty  accumulations  are  saponified  and  become  soluble  in  wa- 
ter. "Drano"  and  other  common  plumbing  cleaners  sold  under  vari- 
ous trademarks  contain  percentages  of  impure  sodium  hydroxide 

Sodium  hydroxide  helps  peel  fruits  and  vegetables.  Many  fruits 
and  vegetables  are  peeled  before  canning  or  dehydration  by  the  use 
of  sodium  hydroxide,  or  lye.  Most  of  the  large  lye  peelers  are  of  the 
moving  conveyor  type,  and  the  fruits  or  vegetables  are  immersed  in  a 
20-  to  25-percent  lye  solution  for  from  two  to  five  minutes.  Durino 
this  time,  the  lye  solution  reacts  with  the  pectins,  or  binding  mate* 
rials,  between  the  individual  cells.  The  skins  become  loose  and  may 
be  removed  by  washing  which  also  removes  all  traces  of  lye. 

This  process  is  an  outgrowth  of  the  making  of  old-fashioned  lye 
hominy,  a  staple  in  the  diet  of  American  pioneers.  In  the  making  of 
lye  hominy,  grains  of  corn  are  soaked  in  a  lye  solution  until  the  hard, 
tough  skin  of  the  corn  grain  becomes  loosened.  Washing  with  fresh 
water  removes  the  skin  and  also  the  lye. 

Other  uses  of  sodium  hydroxide.  Large  quantities  of  sodium  hy- 
droxide are  used  in  reclaiming  aluminum  and  salvaged  rubber,  in 
the  processing  of  many  vegetable  oils,  and  in  the  production  of  gas- 
oline. In  the  refining  of  petroleum,  large  quantities  of  this  "basic 




heavy  chemical"  are  used  to  neutralize  the  sulfuric  acid  with  which 
petroleum  is  purified.  Of  only  slightly  less  importance  is  the  use  of 
sodium  hydroxide  in  digesting  and  purifying  the  cellulose  of  wood 
pulp  that  is  used  in  manufacturing  paper.  Potassium  hydroxide  is 
the  electrolyte  in  certain  types  of  storage  batteries  such  as  the  Edison, 
and  Ni-Cd  batteries.  Lithium  hydroxide  is  used  in  submarines  to 
absorb  CO2. 

How  sodium  hydroxide  is  prepared  for  industrial  use.  Chlorine, 
hydrogen,  and  sodium  hydroxide  are  all  formed  at  the  same  time  dur- 
ing the  electrolysis  of  brine  (see  Chapter  10) .  Both  chlorine  and  so- 
dium hydroxide  are  prepared  industrially  by  this  method.  If  chlo- 
rine is  the  chief  product  desired,  then  sodium  hydroxide  is  the 
byproduct.  That  is,  it  is  a  substance  formed  incidentally  during  the 
preparation  of  another  substance.  If  sodium  hydroxide  is  the  com- 
pound being  manufactured,  then  chlorine  is  the  byproduct. 

The  apparatus  that  is  commonly  used  in  the  industrial  prepara- 
tion of  sodium  hydroxide  by  the  electrolytic  process  is  the  Hooker 
cell.  The  Nelson  diaphragm  cell,  Vorce  cell,  and  mercury  cell  are 
also  used.  The  graphite  anodes  (see  Fig.  44)  are  covered  by  a  so- 
dium chloride  solution  and  are  separated  from  the  cathode  by  an 
asbestos  diaphragm,  which  prevents  the  chlorine  from  mixing  with 
the  sodium  hydroxide.  Chlorine  gas  escapes  through  an  outlet  in  the 
dome  at  the  top  of  the  cell,  and  hydrogen  gas  passes  through  an  out- 
let from  the  steel  screen  cathode.  Because  the  sodium  hydroxide 
solution  is  heavier  than  the  salt  solution,  it  concentrates  with  it  at 
the  bottom  of  the  cell.  It  is  drawn  off,  and  evaporated  to  dryness 
during  which  almost  all  of  the  NaCl  crystallizes  out.  This  process 

Rased  on  a  diagram  by 

Hooker  Electrochemical  Company 




graphite  anode 


Fig.  44.  Hooker  cell  used  in  the  pro- 
duction of  sodium  hydroxide  and 
chlorine.  Hydrogen  is  a  byproduct. 


is  continuous,  and  more  brine  is  added  as  the  strength  of  the  sodium 
chloride  solution  diminishes. 

2NaCl  +  2H2O  -» 2NaOH  +  H2 1  +  C12 1 

An  older  method,  still  very  widely  used,  depends  upon  the  conver- 
sion of  the  cheaper  base,  calcium  hydroxide,  into  sodium  hydroxide 
by  means  of  a  solution  of  sodium  carbonate.  Until  1940,  more  so- 
dium hydroxide  was  produced  by  this  process  than  by  electrolysis. 

Ca(OH)2  +  Na2CO3  -»  CaCO3  [  +  2NaOH 

Since  calcium  carbonate  is  insoluble,  it  is  separated  from  the  soluble 
sodium  hydroxide  by  filtration. 

Methods  of  preparing  a  salt.  Most  inorganic  compounds,  such  as 
sodium  chloride,  sodium  nitrate,  copper  sulfate,  and  so  forth,  are 
salts.  We  have  already  had  occasion  to  refer  to  six  of  the  seven  meth- 
ods of  preparing  salts.  A  list  of  these  seven  methods  follows: 

1)  Neutralization: 

HC1  +  NaOH  -» H2O  +  NaCl 

2)  Action  of  an  acid  on  a  metal: 

3)  Union  of  a  metal  and  a  nonmetal: 

Fe-f  S-+FeS 

4)  Action  of  an  acid  on  the  oxide  of  a  metal: 

2HC1  +  CuO  ->  H2O  +  CuCl2 

5)  Action  of  an  acid  on  a  salt  of  a  more  volatile  acid: 

H2SO4  +  2NaCl  ->  2HC1  +  Na2SO4 

6)  Action  of  one  salt  on  another  salt: 

AgNO3  +  NaCl  ->  AgCl  +  NaNO3 

7)  Action  of  the  oxide  of  a  metal  (basic  oxide)  on  the  oxide  of 
a  nonmetal  (acidic  oxide)  : 

CaO  +  SO8  -»  CaSO4 
The  horizontal  lines  under  certain  formulas  indicate  salts. 



Fabre,  Jean  H.  The  Wonder  Book  of  Chemistry,  pp.  154- 
170.  Albert  &  Charles  Boni,  New  York,  1922.  Salts  and  neu- 
tralization are  discussed  in  a  captivating  manner. 

Jaffe,  Bernard.  Chemical  Calculations,  pp.  96-102.  World 
Book  Co.,  Yonkers,  New  York,  1947.  Normal  and  molar  solu- 
tions and  problems  involving  them  are  included. 


1.  A  base  is  a  substance  that  contains  a  metal  or  metallic 
radical  and  one  or  more  hydroxyl  groups.  Its  water  solution 
is  soapy  to  the  touch,  turns  pink  litmus  blue,  and  reacts  with 
acids,  forming  water  and  a  salt. 

2.  Neutralization  is  the  action  of  an  acid  with  a  base,  form- 
ing water  and  a  salt.  The  hydrogen  of  the  acid  unites  with  the 
hydroxyl  radical  of  the  base,  forming  water. 

3.  When  a  strong  acid  and  a  strong  base  react,  forming  a 
mole  of  water,   13,700  calories  of  heat  are  liberated.  This 
amount  of  heat  is  called  heat  of  neutralization. 

4.  A  salt  is  a  compound  made  up  of  a  metal  or  a  metallic 
radical  and  a  nonmetal  or  an  acid  radical. 

5.  The  seven  methods  of. preparing  salts  ar^:  (1)  neutraliza- 
tion,  (2)  action  of  an  acid  on  a  metal,   (3)  union  of  a  metal 
and  a  nonmetal,    (4)  action  of  an  acid  on  a  metallic  oxide, 
(5)  action  of  an  acid  on  a  salt  of  a  more  volatile  acid,   (6)  re- 
action between  two  salts,  and  (7)  reaction  of  a  basic  oxide  on 
an  acidic  oxide. 


Group  A 

1.  (a)  Write  an  equation  illustrating  the  method  of  pre- 
paring NaOH   in  the  laboratory,    (b)  How  can  the  NaOH 
formed  be  changed  into  a  solid? 

2.  (a)  What  other  base  can  be  made  in  the  same  way? 

(b)  Write  the  equation  for  the  preparation  of  this  other  base. 

(c)  What  is  its  common  name? 

3.  (a)  What  is  a  base?    (b)  List  the  four  properties  by 
which  a  base  can  be  recognized. 

4.  Give  the  names  and  formulas  of  three  bases  other  than 
NaOH  that  are  often  found  at  home. 

5.  (a)  What  is  a  salt?  (b)  What  is  neutralization?  (c)  How 
is  mercerized  cotton  prepared? 



6.  A  soil  is  found  to  be  acid.  To  obtain  the  best  results 
with  certain  crops  a  neutral  soil  is  required.  How  would  a 
farmer  correct  the  acid  condition  of  his  soil? 

0.5M  AlBr3,  and 

7.  Write  four  equations  illustrating  neutralization. 

8.  (a)  What  is  a  molar  solution?  (b)  What  weight  of  acid, 
base,  or  salt  do  the  following  solutions  contain: 

(1)  3M  NaOH,  (2)  0.1M  H2SO4  (3) 
(4)  1.8M  HN03? 

9.  (a)  What  is  a  normal  solution?    (b)  What  weight  of 
acid,  base,  or  salt  do  the  following  solutions  contain: 

(1)  0.1N  Mg(OH)2,    (2)   0.5N  H,PO4,    (3)   2.5N  AgNO3, 
and   (4)    1.5N  A1C13? 

10.  What  is  meant  by  titration? 


11.  Describe    the   electrolytic   process   for   making   NaOH, 
illustrating  your  description  with  a  diagram. 

12.  In  a  process  of  the  kind  mentioned  in  exercise  11,  what 
determines  which  is  the  product  and  which  the  byproduct? 

13.  Write  the  equation  for  a  method  of  preparing  sodium 
hydroxide  other  than  the  electrolytic  method. 

14.  (a)  What  is  the  annual  consumption  of  NaOH  in  the 
United  States?  (b)  What  are  the  three  chief  uses  of  NaOH? 

15.  (a)  What  is  lye?    (b)  Why  is  it  useful  in  the  home? 
(c)  What  cautions  should  be  exercised  in  its  use? 

16.  A  sample  of  fabric  was  placed  in  a  test  tube  containing 
a  solution  of  NaOH,  and  boiled  for  a  few  minutes.  Half  of  the 
fabric  disappeared.  What  can  you  say  about  its  composition? 

1 7.  Copy  and  complete  the  following  table,  inserting  the  cor- 
rect formulas  for  the  salts  formed.  Do  not  write  in  this  book. 










18.  Classify   as   acids,   bases,   or   salts:    A1(OH)8,    K2CO8 
CuSO,,  Pb  (OH)  „  HC2H802,  CaCl2,  BaSO4,  HCN,  H2S. 


19.  Which  will  produce  a  greater  weight  of  alkali,  45  g.  of 
Na  or  50  g.  of  K? 

20.  A  piece  of  Na  is  thrown  in  H2O.  The  NaOH  formed 
exactly  neutralizes  24.5  g.  of  H2SO4.  What  weight  of  Na  was 

Group  B 

21.  (a)  Write  the  equations  for  the  neutralization  of  sodium 
hydroxide  with   (1)  hydrochloric  acid,  and   (2)  sulfuric  acid, 
(b)  The  heat  of  neutralization  for  the  first  reaction  is  13,700 
calories;  for  the  second  reaction,  2  X  13,700  calories.  Inspect 
the  two  equations.  Explain  why  the  heat  of  neutralization  is 
twice  as  large  in  the  second  reaction. 

22.  What  volume  of  a  0.1N   KOH  solution  is  needed  to 
neutralize  50  ml.  of  a  0.5N  solution  of  H2SO4? 

23.  What  is  the  reason  for  adding  lime  water  to  milk  that  is 
to  be  fed  to  an  infant? 

24.  NaHCO3  is  frequently  used  to  neutralize  an  acid,  (a)  Is 
it  a  base?   (b)  Explain  your  answer. 

25.  A  glass   bottle   with   a   glass   stopper   contained   solid 
NaOH.  When  an  attempt  was  made  to  open  the  bottle,  it  was 
found  that  the  glass  stopper  was  firmly  cemented  to  the  neck 
of  the  bottle.  Explain. 

26.  In  making  cream  of  tomato  soup,  a  pinch  of  NaHCO3 
is  added  to  the  tomato  puree  before  the  milk  is  added.  Explain 
the  reason  for  this  practice. 


1.  Take  home  some  pieces  of  litmus  paper  and  test  the  fol- 
lowing substances  for  acidic,  basic,  or  neutral  properties:  vine- 
gar, milk  of  magnesia,  ammonia  water,  table  salt  solution, 
soap  solution,  sugar  solution,  and  the  liquid  in  an  automobile 
battery.  What  do  you  conclude  about  the  importance  of  acids 
and  bases  in  everyday  living? 

2.  With  the  help  of  your  chemistry  teacher,  perform  the  fol- 
lowing experiment:  Pour  an  excess  of  NH4OH  on  a  solution 
of  FeCl3.  You  will  get  a  brown,  sticky  precipitate  of  Fe  (OH)  3. 
Boil  thoroughly,   then  filter.   Wash   the   insoluble   Fe  (OH)  3 
several  times  with  distilled  water.  What  properties  of  soluble 
bases  are  not  possessed  by  this  insoluble  base? 

3.  Using  litmus  paper  as  indicator,  neutralize    (with  con- 
stant stirring)   a  known  volume  of  vinegar  with  some  house- 
hold ammonia.  Report  to  your  class  on  the  relative  strength 
of  these  solutions. 




.  .  .  The  sea  is  the  chemist  that  dis- 
solves the  mountains  and  tlie  rocks, 
pulverizes  old  continents  and  builds 
new,  forever  redistributing  the  solid 
matter  of  the  globe.  Ralph  Waldo 
Emerson,  1803-1882 

Is  chemically  pure  water  found  in  nature?  The  purest  form  of 
water  in  nature  is  rain  water.  Although  it  is  perfectly  safe  for  drink- 
ing purposes,  rain  water  is  not  pure  water,  for  it  is  mixed  with  the 
gases  of  the  atmosphere  and  with  small  amounts  of  dust  and  other 
impurities,  which  it  has  washed  down  from  the  air. 

Eventually  all  water  finds  its  way  to  the  ocean,  the  great  reservoir. 
Rain  may  fall  upon  busy  streets  and  flow  through  sewers  to'  the  sea. 
It  may  fall  on  the  ground  and  furnish  the  water  so  necessary  to  grow- 
ing plants.  Water  evaporates  from  the  leaf-surfaces  of  plants  and 
from  the  surfaces  of  streams,  lakes,  and  oceans  only  to  fall  again  as 
rain.  Without  this  water-cycle,  life  as  we  know  it  could  not  exist. 

Water  that  flows  over  the  ground  (surface  water}  collects  fine 
particles  of  solid  material.  The  size  and  the  weight  of  this  material 
depend  on  the  speed  with  which  the  water  flows,  for  fast-moving  wa- 
ter can  carry  much  heavier  material  than  a  slow-moving  stream. 

Water  that  soaks  into  the  ground  (ground  water]  carries  almost  no 
load  of  this  type,  since  the  soil  acts  as  a  filter  and  holds  the  solid  par- 
ticles back.  But  ground  water  is  still  not  pure  water,  no  matter  how 
clear  and  sparkling  it  may  be,  for  it  contains  minerals  that  have  been 
dissolved  out  of  the  soil. 



What  is  a  solution?  When  a  teaspoonful  of  sugar  is  dissolved  in  a 
glass  of  water,  the  sugar  completely  disappears  from  view,  and  we 
say  that  we  have  made  a  solution  of  sugar  and  water.  The  substance 
that  is  dissolved  is  called  the  solute;  the  substance  in  which  the 
solute  dissolves  is  called  the  solvent.  Thus,  when  we  make  a  solution 
of  sugar  and  water,  the  sugar  is  the  solute  and  the  water  is  the  solvent. 

If  we  taste  a  sample  of  this  solution  of  sugar  and  water,  we  find  that 
it  tastes  the  same,  no  matter  whether  we  have  taken  our  sample  from 
the  top  or  the  bottom  of  the  solution.  This  indicates  that  a  solution 
has  the  same  composition  throughout;  that  is,  it  is  homogeneous. 
If  we  take  care. not  to  let  any  of  the  water  evaporate,  the  sugar  will 
not  settle  to  the  bottom;  the  solution  will  remain  the  same.  Of  course, 
different  amounts  of  sugar  will  dissolve  in  a  glass  of  water  —  the 
solution  may  be  very  sweet  or  it  may  be  only  slightly  sweet.  Thus 
a  solution  differs  from  a  compound,  for  we  can  vary  the  composi- 
tion of  a  solution.  The  composition  of  a  compound  does  not  vary. 

A  uniform  mixture  of  solute  and  solvent  that  does  not  conform  to 
the  law  of  definite  proportions  is  called  a  solution.  The  kinetic 
theory  of  matter  helps  to  explain  the  mechanism  of  solution.  Accord- 
ing to  this  theory,  the  solute  breaks  down  into  molecules  which  dis- 
tribute themselves  between  the  molecules  of  the  solvent.  For  this  rea- 
son, a  solution  is  sometimes  called  a  molecular  dispersion  of  a  solute 
in  a  solvent.  If  the  solute  is  colorless  (for  example,  sugar) ,  it  can  no 
longer  be  identified  by  sight  when  it  is  in  solution. 

Difference  between  dilute  and  concentrated  solutions.  A  pinch  of 
salt  in  a  gallon  of  water  makes  a  very  dilute  solution.  Half  a  pound 
of  salt  added  to  the  same  amount  of  water  makes  a  concentrated  solu- 
tion. When  only  a  small  amount  of  solute  is  dissolved  in  a  solvent,  we 
have  a  dilute  solution;  when  the  amount  of  solute  dissolved  is  con- 
siderable we  have  a  concentrated  solution. 

What  determines  the  amount  of  solute  that  will  dissolve?  There 
is  a  limit,  of  course,  to  the  amount  of  a  solute  that  will  dissolve  in 
a  given  volume  of  a  solvent.  Several  factors  determine  the  amount  of 
solute  that  will  enter  into  solution.  The  most  important  factors  are 
the  nature  of  the  solvent  and  the  nature  of  the  solute.  Large  amounts 
of  certain  substances,  such  as  salt,  dissolve  in  water,  but  the  amount 
of  gold  that  dissolves  in  water  is  extremely  minute.  Iodine  dissolves 
only  slightly  in  water,  but  it  is  very  soluble  in  alcohol,  forming  an 
alcohol  solution  known  as  a  tincture.  Just  why  a  substance  dissolves 
in  one  solvent  and  not  in  another  is  not  thoroughly  understood. 

Temperature  has  a  great  deal  to  do  with  the  amount  of  a  solute 
that  will  enter  into  solution.  More  sugar  will  dissolve  in  hot  tea  than 



in  iced  tea.  In  general,  most  solids  dissolve  in  larger  amount  in  warm 
liquids  than  in  cold  liquids.  This  is  not  true  of  gases,  for  the  higher 
the  temperature  of  a  liquid,  the  smaller  the  amount  of  gas  that  will 
dissolve  in  it.  You  are  familiar  with  this  fact  if  you  have  ever  noticed 
that  gas  from  a  warm  bottle  of  soda  pop  escapes  more  rapidly  than  gas 
from  an  ice-cold  bottle  of  the  same  beverage. 

TABLE  9. 

CURVES         1 3 

Solubility  varies  with  ^ 
I  temperature.  A  nearly  ^  °* 
]  straight  line  indicates   H  § 
I  that  an  increase  in 
I  temperature  produces 
1  a  regular  increase  in 

0      10    20 

30    40    50    60    70 
Temperature,  °C. 

Another  important  factor  that  determines  the  solubility  of  a  gas  is 
pressure.  The  higher  the  pressure,  the  greater  the  amount  of  gas  that 
will  dissolve.  For  this  reason,  when  we  want  to  make  a  concentrated 
solution  of  carbon  dioxide  in  water  (the  carbonated  water  you  see  at 
every  soda  fountain) ,  we  add  the  gas  to  the  water  under  high  pres- 
sure and  at  low  temperature.  The  weight  of  gas  that  dissolves  is  pro- 
portional to  the  pressure  on  the  gas  (Henry's  law) .  In  other  words, 
if  we  double  the  pressure,  twice  as  much  gas  will  dissolve. 

Two  factors  that  help  to  determine  the  amount  of  solute  that  will 
dissolve  in  a  definite  period  of  time  (speed  of  solution)  are  the  degree 
of  subdivision  of  the  solute  and  the  extent  to  which  particles  of  solute 
and  solvent  are  closely  intermingled  by  stirring.  As  you  know,  the 
finer  the  particles  and  the  more  vigorous  the  stirring,  the  quicker  the 
solid  will  dissolve.  However,  while  these  two  factors  affect  the  speed 
at  which  a  solute  dissolves,  they  do  not  affect  the  maximum  amount 
of  solute  that  will  dissolve. 

What  is  a  saturated  solution?  If  we  add  a  salt,  such  as  alum,  to  a 
given  volume  of  water  at  a  definite  temperature  and  under  fixed 

small  crystal 
of  hypo 

solution  of  hypo 

(sodium  thiosulfate) 

of  hypo 


crystals   ' 
starting  to  form 

hypo  crystals 
Fig.  45.  Formation  of  hypo  crystals  from  a  supersaturated  solution 


conditions  of  pressure,  the  water  will  continue  to  dissolve  the  salt  un- 
til it  has  dissolved  a  certain  amount.  After  this  amount  is  dissolved, 
no  more  alum  will  dissolve,  and  any  additional  alum  that  is  added 
will  settle  to  the  bottom  of  the  water  and  remain  there.  A  solution  in 
which  no  more  of  the  solute  will  dissolve,  at  that  particular  tempera- 
ture and  pressure,  is  a  saturated  solution.  As  long  as  the  solvent  will 
dissolve  more  solute,  the  solution  is  said  to  be  unsaturated. 

The  addition  of  another  crystal  of  the  solute  to  a  solution  will 
indicate  saturation  or  unsaturation.  If  the  crystal  added  does  not 
dissolve,  the  solution  is  saturated.  From  the  point  of  view  of  the 
kinetic  theory  of  matter,  some  of  the  crystals  added  to  a  saturated 
solution  do  dissolve,  but  just  as  many  molecules  of  the  solute  come 
out  of  solution.  Thus  the  crystals  added  do  not  appear  to  dissolve. 

How  supersaturation  is  used  in  purifying  solids.  It  may  sound 
contradictory  to  say  that  it  is  possible  to  prepare  a  supersaturated 
solution,  that  is,  a  solution  which  contains  more  of  the  solute  than  a 
saturated  solution.  But  it  is  possible.  For  example,  to  prepare  a  super- 
saturated solution  of  hypo  (Na2S,O8  -  5H,O) ,  we  first  make  a  satu- 
rated solution  of  this  salt  in  boiling  water,  and  then  slowly  cool  the 
solution.  The  excess  hypo  does  not  come  out  of  solution,  as  we  would 
expect,  but  remains  in  solution.  Because  this  solution  contains  more 
solute  than  it  normally  holds  at  the  lower  temperature  when  satu- 
rated, it  is  called  a  supersaturated  solution. 

Supersaturation  is  an  unstable  condition,  and  if  the  solution  is 
disturbed  by  adding  a  tiny  crystal  of  the  solute,  all  of  the  excess  salt 
separates  out  and  the  solution  becomes  saturated.  Since  only  pure 



crystals  will  separate  out,  this  method  is  often  used  commercially  in 
preparing  chemically  pure  (C.P.)  salt  crystals.  Because  impurities 
are  present  to  a  small  degree  as  a  rule,  they  separate  out  later. 

The  effect  of  a  solute  on  freezing  and  boiling  points.  As  you  know, 
the  addition  of  salt  to  water  lowers  the  freezing  point  of  the  solu- 
tion. In  general,  a  solute  raises  the  boiling  point  and  lowers  the 
freezing  point  of  a  solvent  a  definite  number  of  degrees.  This  prin- 
ciple is  used  in  preparing  antifreeze  mixtures,  which  freeze  at  tem- 
peratures much  lower  than  0°C.,  the  freezing  point  of  pure  water. 
A  water  solution  that  is  50  percent  grain  alcohol  does  not  freeze  un- 
til it  reaches  —  31°C.  Hence,  solutions  containing  alcohol  or  other 
solutes  are  used  in  automobile  radiators  in  cold  weather. 

All  substances  do  not  have  the  same  effect  on  freezing  and  boiling 
points.  Certain  acids,  bases,  and  salts  raise  the  boiling  point  and 
lower  the  freezing  point  two  or  three  times  as  much  as  other  sub- 
stances including  sugar,  alcohol,  and  glycerin.  This  singular  behavior 
has  an  important  relation  to  dissociation  (see  Chapter  16) . 

The  difference  between  a  solution  and  a  suspension.  If  a  substance 
does  not  actually  dissolve  but  becomes  mixed  somewhat  uniformly 
with  a  solvent  and  then  separates  on  standing,  the  mixture  is  not 
a  true  solution  but  a  suspension.  Fine  clay  or  other  fine  materials 
intimately  mixed  with  water  is  a  suspension. 

The  chief  differences  in  behavior  between  a  true  solution  and  a 
suspension  depend  upon  differences  in  the  size  of  the  particles  of  the 
solute  in  solution  compared  with  the  size  of  the  particles  in  suspen- 
sion. Particles  of  solutes  in  solution  are  in  the  molecular  state  of 
subdivision.  But  even  the  finest  particles  in  suspension  are  much 
larger  than  molecules;  hence,  they  settle  out. 

Particles  in  suspension  may  be  separated  by  filtration;  those  in  a 
solute  cannot  be  separated  by  filtration.  The  solute  of  a  true  solution 
has  a  greater  effect  upon  the  boiling  and  freezing  points  of  the 
solvent  than  the  particles  in  suspension  have  on  a  suspension. 

In  the  borderland  between  true  solutions  and  suspensions  is  an- 
other condition  of  matter  called  the  colloidal  state.  This  condition 
of  matter  is  so  important  that  all  of  Chapter  38  is  devoted  to  it. 

How  we  purify  water.  Even  before  the  germ  theory  of  disease 
was  proved,  several  methods  were  in  use  for  the  purification  of 
water.  In  ancient  times,  water  was  made  more  fit  to  drink  by  filter- 
ing, boiling,  or  by  allowing  suspended  impurities  to  settle  out.  Laws 
were  passed  to  guard  against  the  contamination  of  river  water  used 
for  drinking  purposes  by  prohibiting  the  washing  of  clothes  and  the 
disposal  of  refuse  and  sewage  in  it. 


In  a  chemical  sense  most  drinking  water  is  not  pure,  tor  it  con- 
tains dissolved  minerals  and  gases.  It  is  pure  in  a  biological  sense, 
which  means  that  it  is  comparatively  tree  from  bacteria  and  other 
organic  matter  harmful  to  health. 

Water  used  for  drinking  purposes  may  be  put  through  one  or 
more  of  the  following  processes  of  purification,  depending  upon  the 
nature  of  the  impurities  it  contains: 

(1)  Aeration  consists  of  spraying  the  water  into  the  air  or  letting 
it  tlow  down  a  series  of  steps  in  thin  sheets  so  that  sunlight  and  oxy- 
gen may  kill  most  bacteria  present.  (2)  Filtration  consists  of  strain- 
ing the  water  through  a  suitable  sieve  (filter) ,  thus  separating  sub- 
stances either  in  suspension  or  afloat.  Sand  filtration  dates  as  far  back 
as  1829,  when  London  purified  its  drinking  water  by  passing  the 
water  through  beds  of  line  sand.  (3)  Chlorination,  discussed  in 
Chapter  10.  (4)  Ozonalion,  discussed  in  Chapter  2.  (5)  Coagulation 
consists  of  adding  alum  or  other  chemicals  that  cause  the  organic 
matter  containing  bacteria  and  other  suspended  material  to  settle 
out.  Coagulation  is  discussed  more  fully  on  page  396.  (6)  Chlor- 
amination,  another  process  of  water  purification  using  both  chlo- 
rine and  ammonia  is  discussed  on  page  253.  (7)  Demineralization 
by  means  of  synthetic  resins  (see  page  191)  . 

Modern  sewage  disposal.  Closely  related  to  the  problem  of  main- 
taining an  adequate  supply  of  safe  drinking  water  is  the  problem  of 
disposing  ol  the  domestic  and  industrial  wastes  that  are  produced 
especially  under  the  crowded  living  conditions  of  modern  cities. 

Sewage,  as  these  wastes  are  called,  includes  chiefly  the  organic 
wastes  that  may  be  disposed  of  by  ft  nl  re  faction,  a  process  that  con- 
sists of  a  combination  of  bacterial  action  and  oxidation.  In  some 
communities,  sewage  is  discharged  directly  into  streams  which  carry 
the  sewage  away.  The  dissolved  oxygen  in  the  water  of  the  stream 
eventually  oxidizes  sewage;  but,  as  a  result,  the  dissolved  oxygen 

Fig.  46.  Water  filtration.  As  the  water  passes 
through  the  filtering  layers,  solid  matter  and 
many  germs  are  removed. 




An  effective  and  relatively 
inexpensive  method  of  puri- 
fying water  is  to  aerate  it 
by  spraying  it  info  the  air. 

is  used  up  and  all  the  fish  and  other  higher  plant  and  animal  life 
normally  found  in  the  stream  are  unable  to  live  in  the  water  because 
they  cannot  obtain  oxygen.  Only  a  few  of  the  lower  forms  of  life, 
such  as  certain  algae  and  bacteria,  can  live  in  the  polluted  stream, 
and  the  water  is  unfit  for  almost  all  purposes.  Oxygen-consuming 
factory  wastes  are  sometimes  handled  in  disposal  wells.  Fortunately, 
the  oxygen  of  the  air  cleanses  the  polluted  stream  in  the  course  of 
its  meanderings.  The  distance  necessary  to  cleanse  the  stream  de- 
pends upon  several  factors  —  the  amount  of  sewage  discharged,  the 
size  of  the  stream,  its  rate  of  flow,  and  so  forth. 

Most  modern  cities  dispose  of  sewage  by  more  modern  and  less 
harmful  methods  of  treatment  such  as  the  activated  sludge  process. 
The  sewage,  which  flows  to  the  modern  treatment  plant  in  great 
quantities  of  flushing  water,  is  first  run  into  closed  tanks  where  the 
solids  settle  as  a  sludge,  or  is  run  through  sieves  and  screens  that 
remove  the  suspended  solids.  Certain  bacteria  present  in  the  sew- 
age decompose  the  sludge,  liberating  both  nitrogen  and  methane 
gas,  and  convert  the  sewage  into  a  nontoxic,  humus-like  waste  that 
may  be  used  as  a  fertilizer.  The  partially  purified  water  may  be 
further  purified  by  one  of  the  methods  of  purification  already  dis- 
cussed, or  it  may  be  discharged  into  a  neighboring  stream  or  body  of 
water.  In  certain  types  of  installations,  the  methane  gas  produced 
during  the  digestion  of  the  sludge  is  burned  in  gas  engines  to  pro- 
duce the  energy  necessary  to  operate  the  sewage-treatment  plant. 

How  water  can  be  made  chemically  pure  by  distillation,  None 
of  the  methods  mentioned  removes  all  impurities  from  water.  Not 
one  of  them  completely  removes  substances  dissolved  in  water.  Dis- 
solved substances  are  completely  removed  from  solution  both  in 
industry  and  in  the  laboratory  by  distillation.  Because  the  boiling 
points  of  substances  differ,  it  is  possible  to  separate  solute  from 



inlet  from 


Fig.  47.  Distillation  of  water  using  a  Liebig  condenser.  Why  should 
cold  water  enter  the  jacket  at  the  lower  end? 

solvent.  In  general,  solids  that  are  dissolved  in  water  have  higher  boil- 
ing points  than  water,  and  remain  behind  after  the  water  has  been 
boiled  off.  Certain  liquids,  including  glycerin,  have  higher  boiling 
points  than  water;  others,  as  alcohol,  have  lower  boiling  points. 
Hence  two  or  more  liquids  in  solution  may  be  separated  by  distilla- 
tion. Gases  in  solution  are  driven  off  soon  after  the  water  is  heated. 

During  distillation,  water  is  first  boiled,  or  evaporated,  and 
then  the  steam,  or  water  vapor,  is  cooled.  This  cooling  condenses  it 
into  water  again.  Distillation  is  thus  a  double  process,  including  both 
evaporation  and  condensation.  The  first  portion  of  the  distillate,  the 
liquid  that  results  from  the  condensation  of  the  vapor,  may  contain 
small  amounts  of  dissolved  gases.  The  final  portion  of  the  distillate 
may  contain  small  amounts  of  liquids  or  even  dissolved  solids  whose 
boiling  points  are  close  to  that  of  water.  If  these  two  portions  are  dis- 
carded, the  rest  of  the  distillate  will  be  free  of  all  impurities. 

Stills  used  for  industrial  purposes  are  made  of  such  material  as 
copper,  steel,  lead-lined  steel,  and  fused  silica  (SiO2) . 

How  water  is  distilled  in  the  laboratory.  The  laboratory  appara- 
tus for  water  distillation  consists  of  a  flask,  a  Liebig  condenser 
made  popular  in  1850  by  Justus  von  Liebig  (ton  le'biK) ,  a  German 
chemist  famed  for  his  contributions  to  organic  and  agricultural 
chemistry,  and  a  receiver.  Impure  water,  which,  for  the  purpose  of 
the  experiment,  may  contain  small  amounts  of  ammonia,  salt, 
and  red  ink,  is  boiled  in  a  sidearm  flask,  as  shown  in  Fig.  47.  A 
thermometer  indicates  the  boiling  point  of  the  solution.  The  steam 
and  water  vapor  pass  into  the  inner  glass  tube  of  the  Liebig  con- 
denser, which  is  surrounded  by  a  glass  jacket  having  a  glass  inlet 
and  outlet  for  water.  To  condense  the  water  vapor  before  it  escapes 
from  the  inner  glass  tube,  cold  water  is  circulated  through  the  outer 
tube  of  the  condenser.  As  a  result  of  this  cooling,  distilled  water 
collects  as  the  distillate. 

Fractional,  partial  vacuum,  and  high  vacuum  distillation.  When 
the  boiling  points  of  the  impurities  are  very  close  to  the  boiling 




point  of  the  solvent,  simple  distillation  is  not  very  effective.  For  ex- 
ample, the  boiling  point  of  grain  alcohol  is  78°C.,  and  if  separated 
from  water  by  a  single  distillation,  the  distillate  is  not  pure. 

Oil  refiners  face  a  difficult  problem  in  separating  the  liquids  that 
make  up  crude  petroleum.  The  boiling  points  of  these  liquids  are 
so  close  together  that  he  must  resort  to  fractional  distillation.  This 
consists  of  heating  the  mixture  carefully  and  separating  it  into  its 
various  liquid  fractions,  which  boil  off  at  different  temperatures.  For 
example,  the  first  fraction  might  boil  off  below  200°C.,  the  next  be- 
tween 200°C.  and  220°C.,  and  so  forth.  In  each  case,  it  is  necessary 
to  purify  each  fraction  further  by  additional  distillations. 

It  often  happens  that  liquids  cannot  be  separated  by  fractional 
distillation,  because  the  high  temperature  required  may  char  or 
otherwise  decompose  the  solute.  In  such  cases  partial  vacuum,  or 
even  high  vacuum  distillation  is  used.  By  reducing  the  pressure  over 
the  solution,  its  boiling  point  is  also  lowered. 

According  to  the  kinetic  theory  of  matter,  the  molecules  of  water 
leave  the  liquid  and  become  water-vapor  molecules  much  more 
easily  when  the  pressure  over  them  is  decreased.  Water,  for  exam- 
ple, under  normal  atmospheric  pressure  boils  at  100°C.  If  this  pres- 
sure is  lowered  to  one-half  normal,  water  will  boil  at  82°C.,  and 
the  process  of  distillation  can  be  carried  through  at  this  lower  tem- 
perature. Thus,  water  is  removed  from  milk  in  making  evaporated 
milk  by  partial  vacuum  distillation.  Water  is  removed  from  sugar- 
cane sap  in  making  granulated  sugar  by  reduced  pressure  distilla- 
tion. High  vacuum  distillation  is  used  in  making  vaccines,  serums, 
antibiotics,  blood  plasma,  frozen  orange  juice,  and  for  coating  metals. 

The  principle  of  distillation  is   used  in   both  the  huge  petroleum 
"cracker"  (left)  and  the  laboratory  fractionating  equipment  (right). 

Shell   Oil    Company 


Distilled  water  has  many  uses.  Because  distilled  water  is  free  from 
dissolved  gases  from  the  air,  it  is  flat  and  insipid  in  taste.  Its  prop- 
erties are  those  of  chemically  pure  water  (see  page  65) .  Ordinarily, 
distilled  water  is  not  used  for  drinking  purposes  on  land.  However, 
at  sea,  drinking  water  is  commonly  prepared  from  sea  water  by  dis- 
tillation. The  economic  production  of  drinking  water  from  sea  wa- 
ter, which  contains  about  3.5  percent  of  dissolved  salts,  is  very  old, 
dating  back  at  least  to  the  time  of  Aristotle  (about  350  B.C.)  . 

Today,  ships  of  nearly  all  kinds  obtain  drinking  water  by  distill- 
ing sea  water  under  reduced  pressure.  Seagoing  vessels  also  produce 
water  for  use  in  boilers  by  distilling  sea  water  under  reduced  pres- 
sure. In  this  way,  all  salts  are  removed,  and  subsequent  treatment  re- 
moves all  traces  of  dissolved  gases.  Consequently,  the  operation  of 
the  boilers  is  not  reduced  in  efficiency  by  the  formation  of  large 
quantities  of  boiler  scale  (see  page  489) . 

Today,  aviators  who  are  forced  down  at  sea,  and  all  abandon-ship 
equipment  of  warships  and  merchant  vessels,  carry  small  kits  that 
may  be  used  to  prepare  safe  drinking  water  from  sea  water  easily 
and  simply.  Such  kits  produce  drinking  water  by  chemical  means. 
Their  chemistry  is  discussed  on  page  491. 

Distilled  water  is  used  in  storage  batteries,  in  which  ordinary  drink- 
ing water  should  not  be  used  because  of  the  harmful  effect  of  its  im- 
purities on  the  plates  of  the  battery.  Distilled  water  is  indispensable 
in  many  scientific  and  industrial  operations.  In  the  preparation  of 
vaccines,  pharmaceuticals,  certain  dyed  textiles,  and  C.P.  chemicals, 
distilled  water  is  commonly  used. 

Heavy  water,  a  powerful  tool  of  research.  The  heavier  hydrogen 
isotopes,  deuterium  and  tritium  (see  page  56) ,  may  be  represented 
respectively  by  the  symbols  D  and  T.  It  has  been  determined  that  the 
nucleus  of  deuterium  contains  one  proton  and  one  neutron.  There- 
fore its  atomic  weight  is  two.  The  nucleus  of  tritium  contains  one 
proton  and  two  neutrons  and  it  has  an  atomic  weight  of  three. 

When  water  is  formed  from  oxygen  and  one  of  the  two  heavy  hy- 
drogen isotopes,  its  molecular  weight  is  greater  than  18,  the  molecu- 
lar weight  of  ordinary  water.  This  water  is  called  heavy  water.  Heavy 
water  can  also  be  formed  from  deuterium  or  tritium  and  any  one 
of  the  three  heavy  isotopes  of  oxygen.  Deuterium  oxide,  D2O,  is 
present  in  ordinary  water  to  the  extent  of  about  one  part  of  D2O 
in  5000  parts  of  water.  Tritium  is  seldom  found  in  nature  but  is  or- 
dinarily made  in  an  atomic  pile. 

Deuterium  oxide  has  been  prepared  in  large  quantities.  It  differs 
to  a  small  extent  from  ordinary  water  in  both  freezing  and  boiling 


points.  Its  maximum  density  occurs  at  11.6°C.  rather  than  at  4°C., 
as  is  the  case  with  H2O.  Heavy  water  is  used  in  atomic  piles  as  a 

Deuterium  oxide  appears  to  arrest  the  growth  of  seedlings;  tad- 
poles die  prematurely  in  it;  it  is,  however,  not  toxic  to  man  and  the 
higher  animals. 

Heavy  hydrogen  has  been  substituted  for  ordinary  hydrogen  in 
certain  fats,  and  the  course  and  changes  which  these  "tagged"  fat 
molecules  have  undergone  on  their  way  through  the  animal  body 
have  been  studied  by  the  tracer  technique. 


Cerna,  Wendell  W.  "Industrial  Water  Conditioning  Proc- 
esses/' Journal  of  Chemical  Education,  March,  1943,  pp.  107- 
115,  and  April,  1943,  pp.  191-197. 

Ellis,  Cecil  B.  Fresh  Water  from  the  Ocean.  Ronald  Press 
Company,  New  York,  1954.  This  is  a  conservation  study  deal- 
ing with  water  for  cities,  industry,  and  irrigation. 

Goldblatt,  L.  A.,  Ed.  Collateral  Readings  in  Inorganic  Chem- 
istry. D.  Appleton-Century  Co.,  New  York,  1937  (2nd  series, 
1942).  No.  8  of  the  31  articles  in  this  collection  deals  with 
"Factors  Contributing  to  Quality  of  Public  Water  Supplies." 
Written  by  H.  E.  Jordan. 


1.  A  solution  is  a  uniform  mixture  of  solvent  and  solute 
which  does  not  conform  to  the  law  of  definite  proportions. 

2.  A  dilute  solution  contains  very  little  solute  in  comparison 
with   the  solvent;   a  concentrated  solution  contains  a  large 
amount  of  solute. 

3.  Some  of  the  factors  which  determine  the  amount  of  solute 
that  will  dissolve  in  a  solvent  are:    (1)  the  nature  of  solute 
and  solvent,   (2)  temperature,  and  (3)  pressure.  The  speed  of 
solution  depends  upon    (1)  the  state  of  subdivision  of  the 
solute,  and    (2)  how  thoroughly  the  solute  and  solvent  are 

4.  All  solutes  raise  the  boiling  point  and  lower  the  freezing 
point  of  the  solvent.  The  amount  of  solute  determines  the 
number  of  degrees  of  change.  Acids,  bases,  and  salts  in  solution 
affect  boiling  and  freezing  points  to  a  greater  degree  than  do 
other  substances,  including  alcohol,  sugar,  and  glycerin. 


5.  Solutions  and  suspensions  differ  in  the  following  respects: 

(1)  In  a  suspension,  the  mixture  separates  on  long  standing. 

(2)  The   particles   in   suspension   are   much   larger   than   the 
particles  of  a  solute  in  a  solvent   (which  are  in  the  molecular 
state  of  subdivision) .    (3)  The  particles  of  a  suspension  may 
be  separated  out  by  filtration;  the  particles  of  a  solution  can- 
not be  so  separated.    (4)  A  solute  has  greater  effect  on  the 
freezing  and  boiling  points  than  has  a  material  in  suspension. 

6.  The  colloidal  state  of  matter  is  in  the  borderland  be- 
tween true  solutions  and  suspensions. 

7.  Distillation   consists   of  evaporation   and  condensation. 
Impurities  can  be  removed  by  distillation  because  of  the  dif- 
ference in  the  boiling  points  of  a  solvent  and  its  solute. 


Group  A 

1.  (a)  What  is  the  purest  form  of  H2O  found  in  nature? 
(b)  What  property  of  H2O  makes  it  almost  impossible  to  find 
pure  H2O  in  nature? 

2.  (a)  What  is  a  solution?    (b)  In  a  solution  of  NaCl  in 
H2O,  which  is  the  solvent  and  which  the  solute?   (c)  What  is 
a  tincture? 

3.  Distinguish  a  dilute  solution  from  a  concentrated  so- 

4.  Name  four  factors  that  determine  the  degree  and  the 
speed  of  solubility  of  a  substance. 

5.  Why  is  more  sugar  or  more  stirring  required  to  sweeten 
iced  coffee  than  hot  coffee? 

6.  Why  should  sealed  bottles  of  H2O  heavily  charged  with 
CO2  be  kept  cold? 


7.  How  would  you  proceed  to  prepare  a  saturated  water 
solution  of  washing  soda,  Na2CO3? 

8.  How  could  you  change  a  saturated  solution  of  CuSO4  to 
an  unsaturated  solution  having  the  same  volume? 

9.  Without  changing  the  temperature  or  volume  of  a  solu- 
tion of  CO2  in  H2O,  how  could  you  change  an  unsaturated 
solution  of  this  gas  into  a  saturated  one? 

10.  How  could  you  determine  whether  a  solution  is  satu- 
rated, unsaturated,  or  supersaturated? 

11.  Mention  one  commercial  use  of  supersaturation. 


12.  What  is  the  effect  of  a  solute  on  the  freezing  and  boiling 
points  of  a  solvent? 

13.  What  principle  is  involved  in  the  use  of  an  antifreeze 
mixture  in  an  automobile  radiator? 

14.  How  do  true  solutions  and  suspensions  differ? 

15.  What  is  the  name  given  to  the  state  of  matter  that  is  on 
the  borderland  between  true  solutions  and  suspensions? 

16.  What  are  seven  methods  used  to  purify  drinking  water? 

17.  How  is  sewage  disposed  of  in  modern  sewage-disposal 

18.  What  may  be  the  results  of  ineffective  methods  of  sewage 

.  .  I .  . . 


19.  Which  method  or  methods  of  purifying  H2O  produce 
chemically  pure  H2O? 

20.  (a)  What  is  distillation?    (b)  Upon  what  fact  does  the 
separation  of  impurities  by  distillation  depend? 

21.  Why  is  the  middle  portion  of  a  distillate  the  purest? 

22.  Make  a  diagram  of  the  apparatus  used  in  the  laboratory 
distillation  of  H2O. 

23.  (a)  Is  the  drinking  water  of  a  large  city,  such  as  Chicago 
or  New  York,  distilled?  (b)  Give  reasons  for  your  answer. 

24.  Petroleum  is  refined  by  fractional  distillation.  Why? 

25.  Why  are  liquids  often  distilled  under  reduced  pressure? 

26.  State  the  physical  properties  of  distilled  water. 

27.  For  what  is  distilled  water  used? 

28.  What  are  the  desirable  characteristics  of  drinking  water? 

29.  What  is  the  difference  between  wholesome  water  and 
chemically  pure  water! 

30.  (a)  What  is  heavy  water?   (b)  How  does  it  differ  from 
ordinary  water?   (c)  Name  two  uses  of  heavy  water. 

Group  B 

31.  Devise  an  experiment  to  show  that  perfectly  clear  spring 
water  contains  impurities. 

32.  Small  filters  attached  to  household  faucets  sometimes 
become  a  menace  to  health.  Explain  this  statement. 


33.  What  would  be  the  effect  of  the  continued  use  of  rain 
water  in  storage  batteries? 

34.  Is  water  obtained  by  melting  ice  from  a  lake  purer  than 
the  water  of  that  lake?  Explain  your  answer. 

35.  If  a  liquid  is  colorless,  odorless,  and  clear,  how  could 
you  determine  whether  it  is  a  solution  or  a  pure  compound? 

36.  How  does  a  solution  of  NaCl  in  water  differ  from  a  mix- 
ture of  NaCl  and  sugar? 

37.  Can  a  dilute  solution  be  a  saturated  solution  also?  Ex- 

38.  Explain  the  operation  and  principle  of  the  pressure 

39.  (a)  Explain  solution  by  means  of  the  kinetic  theory  of 
matter,   (b)  According  to  the  kinetic  theory,  explain  why  it  is 
easier  to  evaporate  or  distill  a  liquid  under  reduced  pressure. 

40.  Would    there   be    any   advantage   or   disadvantage    in 
making  solutions  of    (a)    table  salt,    (b)    calcium  hydroxide, 
with  hot  water  instead  of  cold?  (c)  Explain  each  answer. 


1.  Visit  your  local  sewage-disposal  plant,  and  study  its  op- 
eration. Make  a  report  on  this  visit  to  your  class. 

2.  Obtain  some  Rochelle  salt    (U.S.P.)    from  your  teacher 
or  druggist  and  prepare  pure  crystals  from  it.  Bring  them  to 
class  and  explain  the  process  you  used.   (As  a  substitute,  you 
may  prepare  "rock  candy"  crystals.) 

3.  Harold  C.  Urey,  discoverer  of  heavy  water,  in  a  lecture  in 
1938,  made  the  following  statement,  "I  believe  I  speak  for  the 
vast  majority  of  all  scientific  men.  Our  object  is  not  to  make 
jobs  and  dividends.  These  are  a  means  to  an  end,  merely  inci- 
dental. We  wish  to  abolish  drudgery,  discomfort  and  want 
from   the  lives  of  men,  and  bring  them   pleasure,   comiort, 
leisure  and  beauty.  Often  we  are  thwarted  but  in  the  end  we 
will  succeed."  Write  a  short  report  either  for  or  against  this 
view,  or  organize  a  class  discussion  on  this  subject. 

4.  Study  the  methods  of  water  purification  in  your  own 
community.  Report  to  your  class  on  this  subject.  Find  out,  if 
you  can,  the  mineral  content  of  the  water  and  the  amount  and 
kind  of  suspended  solids  present  both  before  and  after  treat- 

5.  Using  hypo    (sodium  thiosulfate)    prepare  10  ml.  of  a 
supersaturated  solution.  Allow  the  solution  to  cool  slowly  to 
room  temperature.  Then  add  a  crystal  of  hypo.  Repeat  this 
before  your  class  if  it  can  be  arranged. 



.  .  .  /  heard  Cleve  say:  "Do  you 
believe  sodium  chloride  is  dissolved 
into  sodium  and  chlorine?  In  this 
glass  I  have  a  solution  of  sodium 
chloride.  Do  you  believe  there  are 
sodium  and  chlorine  in  it?"  "Oh, 
yes,"  Ostwald  replied,  "there  is  some 
truth  in  that  idea."  .  .  .  Cleve  threw 
a  look  at  Ostwald  which  clearly 
showed  that  he  did  not  think  much 
of  his  knowledge  of  chemistry. 

Svante  Arrhenius,  1925 

Two  eternal  questions:  How?  and  Why?  Science  is  constantly  try- 
ing to  answer  two  questions  —  how  and  why.  Often  it  is  not  too  dif- 
ficult to  answer  the  hows  but  the  whys  —  well,  that  is  a  different 
story.  Theories  must  be  formulated,  tested,  and  adapted  to  keep 
them  in  accord  with  all  the  observed  facts,  and  that  is  a  big  job. 

At  least  two  great  theories  underlie  much  of  chemistry:  the  elec- 
tron theory  and  the  theory  of  dissociation.  In  a  sense,  the  theory  of 
dissociation  is  but  an  aspect  of  the  electron  theory,  for  the  theory 
of  dissociation  is  explained  in  terms  of  the  electron  theory.  These 
two  great  theories  explain  some  of  the  hows  and  whys  of  chemistry. 

A  Swedish  boy  tackles  some  puzzling  questions.  Why  is  distilled 
water  a  nonconductor  of  electricity?  Why  do  water  solutions  of  some 
substances  conduct  electricity  yet  water  solutions  of  other  substances 
do  not?  What  causes  some  acids  and  bases  to  be  strong  while  others 
are  weak? 

These  were  some  of  the  unsolved  problems  that  confronted 
chemists  when  Svante  Arrhenius  (ar-ra'ni-us)  was  still  at  school  in 

Arrhenius  not  only  wondered  about  these  problems  but  set  to  work 
to  solve  them.  He  had  some  unusual  notions  of  his  own  about  the 



way  in  which  electricity  passes  through  solutions.  Day  after  day,  and 
often  far  into  the  night,  he  worked  in  his  laboratory  with  hundreds 
of  different  solutions.  For  two  years  he  toiled  ceaselessly. 

Arrhenius  attempted  to  formulate  a  theory  that  would  explain 
what  he  had  observed.  In  those  days  his  whole  world,  both  of  wak- 
ing and  sleeping  hours,  was  one  of  solutions,  electric  currents,  and 
formulas.  For  him  the  rest  of  the  world  did  not  exist.  One  night  he 
sat  up  very  late.  Suddenly,  like  a  flash,  he  saw  the  answer  to  the  great 
riddle.  "I  got  the  idea/'  he  wrote,  "in  the  night  of  the  17th  of  May 
in  the  year  1883,  and  I  could  not  sleep  that  night  until  I  had  worked 
through  the  whole  problem." 

He  went  to  his  teacher  of  chemistry.  "I  have  a  new  theory  of 
electrical  conductivity/'  Arrhenius  told  him.  The  professor  looked 
at  this  boy  and  said:  "You  have  a  new  theory?  That  is  very  interest- 
ing. Good-by."  But  Arrhenius  did  not  lose  heart.  He  wrote  to  the 
leaders  in  chemistry.  Most  of  them  were  hostile  to  his  revolutionary 
theory.  After  a  long  struggle,  however,  they  were  forced  to  admit  its 
probable  truth  and  later  saw  Arrhenius  awarded  the  Nobel  prize  in 
chemistry.  And  so  it  happened  that  a  mere  boy,  with  the  clear  insight 
and  the  creative  imagination  of  a  truly  great  scientist,  stepped  in  and 
cleared  away  an  obstacle  that  had  stood  squarely  in  the  path  of  chem- 
ical progress. 

How  Arrhenius  explained  the  conductivity  of  solutions.  The  the- 
ory proposed  by  Arrhenius  is  known  as  the  theory  of  ionization,  or 
dissociation.  Arrhenius  assumed  that  when  an  electrolyte,  such  as 
sodium  chloride,  dissolves  in  water,  it  tends  to  dissociate,  or  ionize. 
That  is,  it  tends  to  break  apart  into  electrically  charged  atoms  or 
groups  of  atoms  (radicals) .  Arrhenius  used  the  term  ion,  meaning 
wanderer,  to  refer  to  an  atom  or  group  of  atoms  carrying  an  electric 
charge.  He  represented  the  dissociation  of  the  sodium  chloride  mole- 
cule into  ions  when  it  dissolves  in  water  thus: 

NaCl  — »  Na+  (sodium  ion)  +  Cl~  (chloride  ion) 
He  represented  the  dissociation  of  the  sodium  nitrate  molecule  thus: 


Arrhenius  could  not  see  the  ions  in  solution.  They  are  far  too  small 
to  be  seen.  He  advanced  the  theory  that  they  were  present  because 
he  could  account  for  what  he  observed  only  by  assuming  such  ions  to 
be  present.  To  him,  when  an  electrolyte  dissolves,  a  certain  number 
of  its  molecules  immediately  split  up  into  ions.  Thus  an  electrically 


neutral  compound  tends  to  dissociate  into  a  number  of  positively  and 
negatively  charged  particles,  or  ions.  They  move  about  in  all  direc- 
tions until  the  passage  of  an  electric  current  draws  each  ion  to  the 
electrode  bearing  an  opposite  charge. 

It  is  the  ions  that  carry  an  electric  current,  or  flow  of  electrons. 
Hence  a  substance  that  does  not  dissociate  into  ions  is  a  nonelectro- 
lyte,  and  a  substance  of  which  a  large  proportion  dissociates  in  water 
solution  is  a  good  electrolyte.  Arrhenius'  theory  of  ions  and  of  elec- 
trolytic dissociation  is  today,  with  certain  modifications,  universally 
accepted  as  the  correct  explanation  of  the  conductivity  of  solutions. 

How  an  ion  differs  from  an  atom.  There  are  two  main  differences 
between  an  atom  and  an  ion:  (1)  An  atom  is  electrically  neutral;  an 
ion  is  positively  or  negatively  charged.  (2)  An  atom  always  consists 
of  a  single  element;  an  ion  may  consist  of  one  or  more  than  one  ele- 
ment, as  in  the  case  of  the  ammonium  ion,  (NH4)  +.  A  sodium  atom 
is  quite  different  from  a  sodium  ion.  The  former  is  a  silvery  metal 
particle  that  reacts  violently  with  water;  the  latter  is  a  colorless  par- 
ticle that  has  no  noticeable  reaction  with  water. 

The  dissociation  of  zinc  chloride  may  be  represented  thus: 

Note  the  change  from  C12  in  ZnCl2  to  2  Cl~  when  the  zinc  chloride 
is  dissociated  into  ions.  The  number  of  positive  or  negative  charges 
on  the  ion  is  equal  to  the  valence  of  the  element  or  radical. 

showing  what  happens  when  the  following  electrolytes  dis- 

a)  Nitric  acid,  HNOr  g)  Sulfuric  acid,  H2SO4. 

b)  Lithium  hydroxide,  LiOH.     h)  Potassium  carbonate, 

c)  Potassium  hydroxide,  K2CO3. 

KOH.  i)  Sodium  phosphate, 

d)  Ammonium  hydroxide,  Na3PO4. 

NH4OH.  j)  Potassium  chlorate,  KC1O3. 

e)  Barium  hydroxide,  k)  Magnesium  bicarbonate, 
Ba(OH)2.  Mg(HC08)2. 

f)  Aluminum  chloride,  A1C13.      1)  Calcium  nitrate, 


Where  and  how  are  free  ions  formed?  Most  commonly,  dissocia- 
tion takes  place  in  a  water  solution.  However,  it  also  occurs  in  com- 
pounds such  as  sodium  hydroxide,  NaOH,  when  they  are  heated  un- 
til they  melt,  or  fuse.  Such  compounds  are  made  up,  not  of  molecules, 
as  was  previously  supposed,  but  of  ions  in  a  definite  pattern  (see 


direction  of  flow 
of  electrons 

heated  filam 

Fig.  48.  Simplified  diagram  of  a  radio 
tube  showing  the  hot  filament  emitting 
a  stream  of  electrons. 

page  631).  As  the  compound  is  heated,  its  ions  vibrate  more  rap- 
idly and  eventually  tear  apart,  leaving  them  free  to  carry  an  electric 
current.  Gases,  too,  at  very  high  temperatures  dissociate  into  ions, 
thus  becoming  conductors  of  electricity.  Flames,  x-rays,  and  radio- 
active elements  ionize  the  air  around  them. 

A  metal,  when  heated  to  incandescence,  boils  off  electrons  from 
the  outer  rings  of  its  atoms.  This  principle  is  used  in  the  electron 
tubes  so  familiar  to  us  in  radios,  television  sets,  and  other  electronic 
instruments.  Certain  types  of  electron  tubes  are  filled  with  a  gas  or 
mixture  of  gases  at  low  pressure.  Electrons  from  the  hot  cathode  of 
the  tube  strike  the  molecules  of  the  gas,  ionizing  them. 

The  positive  ions  move  toward  the  negative  electrode,  or  cath- 
ode, and  the  negative  ions  and  electrons  move  toward  the  positive 
electrode,  or  anode.  In  a  radio  tube,  the  anode  is  the  plate,  and 
the  cathode  may  be  the  hot  filament.  Thus  a  stream  of  electrons  or 
electric  current  flows  through  the  tube.  Because  a  gas-filled  tube 
contains  more  molecules  that  can  be  ionized  than  does  a  vacuum 
tube  from  which  most  of  the  molecules  have  been  removed,  relatively 
large  currents  can  flow  through  a  gas-filled  tube.  The  thyratron  tube, 
often  employed  as  a  rectifier,  is  a  common  example  of  a  gas-filled 
electron  tube. 

How  Arrhenius  explained  the  action  of  acids.  All  acids  in  water 
solution  contain  free  hydrogen  ions.  These  hydrogen  ions  determine 
the  typical  properties  of  acids.  It  is  the  free  hydrogen  ion  that  turns 
blue  litmus  pink,  has  a  sour  taste,  is  replaceable  by  a  metal,  and 
neutralizes  a  base  by  combining  with  its  hydroxyl  ion. 

The  number  of  hydrogen  ions  present  determines  the  strength  of 
an  acid.  A  strong  acid  is  one  that  dissociates  easily,  and  thus  pro- 
duces a  large  number  of  hydrogen  ions.  The  most  commonly  used 
strong  acids  are  hydrochloric,  nitric,  and  sulfuric  acids.  Carbonic 
acid,  H2CO3,  sulfurous  acid,  H2SO3,  and  boric  acid,  H3BO8,  are  weak 
acids,  because  they  do  not  dissociate  easily,  and  thus  form  only  a 
small  number  of  hydrogen  ions  in  solution. 


How  Arrhenius  explained  the  action  of  bases.  What  has  been 
said  about  strong  and  weak  acids  refers  equally  well  to  strong  and 
weak  bases,  except  that  the  determining  factor  here  is  the  number  of 
free  hydroxyl  (OH)  ~  ions  present.  In  water  solution,  a  strong  base, 
such  as  sodium  hydroxide,  NaOH,  potassium  hydroxide,  KOH,  or 
lithium  hydroxide,  LiOH,  forms  large  numbers  of  hydroxyl  ions. 
A  weak  base,  such  as  ammonium  hydroxide,  NH4OH,  produces  a 
comparatively  small  number  of  (OH)-  ions. 

Some  bases,  like  copper  hydroxide,  Cu  (OH)  ,,  and  aluminum  hy- 
droxide, Al  (OH)  3,  are  extremely  weak,  because,  in  addition  to  the 
fact  that  they  form  a  relatively  small  number  of  ions,  they  are  also 
only  sparingly  soluble.  The  properties  of  bases  are  caused  by  the  pres- 
ence of  free  hydroxyl  ions.  Hence  dry,  solid  sodium  hydroxide  has  no 
basic  properties;  that  is,  it  does  not  react  as  a  base. 

How  the  theory  of  Arrhenius  explained  neutralization.  The  hy- 
drogen ion  and  the  hydroxyl  ion  are  responsible  for  acidic  and  basic 
properties  respectively.  If  these  free  ions  are  removed,  acidic  and 
basic  properties  are  destroyed.  This  is  exactly  what  happens  in  neu- 
tralization. The  hydrogen  ion  of  the  acid  unites  with  the  hydroxyl 
ion  of  the  base,  forming  water,  ;i  nonelectrolyte. 

Na+  +  (OH)~  +  H+  +  Cl~  ->  Na+  +  Cl~  +  HOH 

Hydrogen-ion  concentration  (pH)  and  how  it  is  measured.  It  is 

frequently  important  to  know  whether  complete  neutralization  has 
been  produced,  or  what  the  degree  of  acidity  or  alkalinity  (basic- 
ity) of  a  substance  is.  Sugar-refiners,  brewers,  paper-makers,  electro- 
platers,  sanitary  engineers,  and  bacteriologists  must  have  a  working 
knowledge  of  the  acidity  or  alkalinity  of  many  substances. 

Chemists  make  use  of  a  hydrogen-ion  scale  in  which  the  unit  is 
the  /;H  value  (pronounced  by  reading  the  letters  p,  H)  ,  just  as  the 
unit  of  temperature  is  a  degree.  A  jm  value  of  seven  is  considered 

Chemists  in  professional  labora- 
tories use  special  equipment  to 
determine  the  pH  value  of  sol- 



true  neutralization.  Pure  water  has  such  a  value.  A  pn  value  less  than 
seven  indicates  an  acid  condition.  Thus,  saliva  has  a  pn  value  of 
about  6.9  and  is  slightly  acid.  A  pn  value  greater  than  seven  indi- 
cates an  alkaline  condition.  Thus,  normal  blood  has  a  pu  value  of 
about  7.3  and  is  slightly  alkaline.  A  value  higher  than  7.5  repre- 
sents a  condition  of  alkalosis.  Since  the  human  blood  is  normally 
slightly  alkaline,  a  pn  below  7.2  indicates  a  condition  of  acidosis. 
A  pn  value  of  five  represents  an  acidity  10  times  as  great  as  a  pn 
of  six. 

Litmus  may  be  used  as  an  indicator  only  with  those  solutions 
whose  pn  is  not  less  than  four  or  greater  than  eight.  Other  indi- 
cators have  wider  or  narrower  ranges.  By  comparison  with  standard 
color  tubes  for  each  indicator,  the  degree  of  acidity  or  alkalinity 
of  a  liquid  can  be  determined  accurately. 

Several  types  of  indicator  papers  have  been  developed  for  deter- 
mining the  pn  values  of  solutions.  The  use  of  such  papers  contain- 
ing water-soluble  dyes  is  very  convenient  and  is  much  quicker  than 
referring  to  standard  color  tubes.  One  such  paper  is  known  as 
"Hydrion  Paper." 

How  the  degree  of  dissociation  was  determined.  We  have  seen 
that  some  substances  produce  a  large  percentage  of  ions  and  other 
substances  produce  a  small  percentage  of  ions.  In  other  words,  com- 
pounds differ  in  the  extent  to  which  they  form  ions  in  solution,  that 
is,  in  their  degree  of  dissociation.  The  degree  of  dissociation  of  a 
solution  depends  upon  (7)  the  solute,  (2)  the  solvent,  (3)  the  con- 
centration of  the  solution,  and  (4)  the  temperature. 

The  degree  of  dissociation  is  measured  by  electric  conductivity, 
that  is,  the  ease  with  which  an  electric  current  passes  through  a 
solution.  This  may  be  determined  roughly  by  setting  up  an  appa- 
ratus as  shown  in  the  illustration  below  (Fig.  49)  .  When  the  bottle 
contains  carbonic  acid  (a  weak  acid)  the  electric-light  bulb  glows 
faintly.  When  it  contains  a  dilute  solution  of  hydrochloric  acid  (a 
strong  acid)  ,  the  light  bulb  glows  brightly,  showing  that  little  resist- 
ance is  being  offered  to  the  passage  of  the  current.  A  solution  of 
sugar  in  water,  when  placed  in  the  bottle,  does  not  produce  a  glow 
in  the  bulb,  and  thus  shows  that  a  sugar  solution  is  a  nonelectrolyte. 

Fig.  49.  Laboratory  setup  for  studying 
the  conductivity  of  a  liquid.  When  will 
the  bulb  light? 

platinum  electrodes 
liquid  to  be  tested 


Arrhenius  determined  the  apparent  percentage  of  dissociation  of 
many  compounds  by  this  method.  He  found  that  in  one-tenth  nor- 
mal water  solution  (0.1N)  the  apparent  dissociation  of  hydrochloric 
acid,  nitric  acid,  and  sodium  hydroxide  was  about  92  percent  (strong 
electrolytes)  ;  that  of  sulfuric  acid  about  61  percent;  that  of  potas- 
sium chloride  about  86  percent;  and  that  of  acetic  acid,  ammonium 
hydroxide,  and  mercuric  chloride  only  about  1  percent  (weak  elec- 
trolytes) .  Later  it  was  found  that  this  method  had  certain  limitations. 

How  Arrhenius  explained  abnormal  boiling  and  freezing  points 
of  solutions.  The  theory  of  ionization  also  explains  why  certain 
ac  ids,  bases,  and  salts  raise  the  boiling  points  of  their  water  solu- 
tions to  an  abnormally  high  degree,  though  sugar  does  not  to  the 
same  extent.  The  acid  dissociates  and  produces  two  or  three  times 
as  many  particles  (ions)  as  there  are  molecules  of  undissociated  sugar. 
The  higher  percentage  of  dissociation  of  the  acid  produces  a  greater 
increase  in  the  boiling  point,  since  it  is  the  actual  number  of  par- 
ticles (ions  or  molecules)  in  solution  that  determines  both  the  boil- 
ing and  freezing  points  of  a  solution. 

How  the  theory  of  ionization  fits  in  with  the  electron  theory. 
When  Arrhenius  proposed  his  theory  of  dissociation,  the  electron  was 
still  undiscovered,  and  the  electron  theory  of  matter  had  not  yet 
been  formulated.  Ions,  however,  fit  in  beautifully  with  our  present 
electron  concept  of  matter.  Atoms  become  ions  by  gaining  or  losing 
electrons.  For  example,  if  an  electrically  neutral  atom  of  sodium 
(Na±)  loses  1  electron,  it  becomes  a  positively  charged  particle,  and 
we  represent  it  Na+.  This  is  a  sodium  ion  (Na°  —  €  — >  Na+)  .  Sim- 
ilarly, if  an  electrically  neutral  atom  of  chlorine  (Cldb)  gains  I  elec- 
tron, it  becomes  a  negatively  charged  particle,  and  we  write  it  Cl~. 
This  is  a  chloride  ion  (Cl°  +  €  ->  Cl~)  . 

The  compound  made  up  of  these  elements  (NaCl)  is  really  com- 
posed of  both  of  these  ions  held  strongly  together,  making  what  is 
known  as  an  ion-molecule  that  is  electrically  neutral.  However,  when 
these  ion-molecules  are  placed  in  water,  they  are  split  into  two  parts, 
and  so  long  as  water  is  present,  we  have  separate  sodium  ions  and 
chloride  ions.  Dissociation,  then,  may  be  represented  as  shown  below. 

Fig.  50.  Dissociation  of  a  sodium  chloride  molecule,  ac- 
cording to  the  electron  thepry. 

NaC1  -molecule  *-      Na+ (ion)      +      CT  (ion) 

(+11-11-0)  (-H7-17-0)  (+11-10-+1)       (+17-18--1) 


The  charge  on  the  ion  is  thus  seen  to  be  the  same  as  the  valence  of 
the  element. 

How  the  theory  of  Arrhenius  was  later  modified.  In  general, 
Arrhenius'  theory  has  stood  the  test  of  time  very  well.  In  three  re- 
spects, however,  it  has  been  modified  slightly. 

1)  The  Swedish  chemist  thought  that  water  simply  kept  the  ions 
apart,  but  today  we  have  a  better  understanding  of  the  impor- 
tant role  of  the  solvent.  The  covalent  water  molecule  is  pictured 
as  a  tiny  magnet  with  a  plus  hydrogen  end  and  a  less  positive  oxy- 
gen end.  Such  a  molecule  is  called  a  dipole. 

When  crystals  of  the  polar  compound,  NaCl  (made  up  of  a 
lattice  of  Na  and  Cl  ions  —  see  page  630) ,  are  added  to  water,  the 
positive  end  of  the  dipole  molecule  attracts  the  negative  Cl  ion; 
the  other  end  of  the  dipole  attracts  the  positive  Na  ion.  The  Na 
and  Cl  are  thus  dissociated  and  free  sodium  and  chloride  ions  are 
produced  in  solution.  When  the  water  is  evaporated,  the  ions  re- 
combine  forming  ion-molecules  of  NaCl  crystals.  NaCl  and  other 
ionic  compounds  are  strong  electrolytes. 

2)  The  Br0nsted-Lowry  theory  has  added  to  what  Arrhenius  be- 
lieved about  the  ionization  of  acids.  According  to  this  theory,  an 
acid  is  a  proton    (hydrogen  ion)    donor;  that  is,  any  compound 
that  tends  to  lose  a  proton  (H+)   to  another  substance  is  an  acid. 

For  example,  in  the  case  of  a  water  solution  of  HC1,  the  water 
molecule  combines  with  the  positive  hydrogen  ion  end  (proton) 
of  the  HC1,  forming  a  hydrogen  (hydronium)  ion,  leaving  the 
chlorine  negatively  charged,  thus: 

H2  O  +  H      Cl  ^H20  .  H+  +   Cl" 

V^^CI)         hydronium  ion 

As  HC1  dissolves  in  water,  heat  is  liberated,  indicating  a  chemical 
change  is  taking  place.  The  hydronium  ion  is  an  acid  since  it  may 
give  up  its  proton.  It  is  also  written  as  H8O+. 

A  base  is  a  proton  acceptor.  A  base  is  any  substance  which  com- 
bines with  a  hydrogen  ion  or  proton.  Thus  the  water  molecule 
which  will  accept  a  proton  is  a  base,  though  not  an  active  one. 
Bases  are  often  negative  ions  such  as  OH~,  which  is  an  active  one. 

3)  Insofar  as  weak  electrolytes  are  concerned,   the  theory  ex- 
pressed by  Arrhenius  is  still  correct.  They  are  not  completely  dis- 
sociated. In  the  case  of  strong  electrolytes,  however,  it  is  now  be- 
lieved that  they  dissociate  completely.  The  fact  that  the  heat  of 


hydrogen^  ^JB^Wiiii^^Xoxygen 

Fig.   51.   Action  during  the 
electrolysis  of  water. 

neutralization  of  all  strong  acids  and  bases  is  the  same  is  one  evi- 
dence in  support  of  this  belief.  The  fact  that  when  dilute  solu- 
tions of  strong  electrolytes  are  mixed  there  is  no  trace  of  heat  is 
further  proof.  The  degree  of  dissociation  of  weak  electrolytes  can 
be  determined  accurately,  but  that  of  strong  electrolytes  cannot 
because  of  the  disturbing  electrical  effects  of  a  large  concentration 
of  ions.  In   such  a  high  concentration,  ions  do  not  behave  as  "free 
ions"  and  only  an  apparent  ionization  is  obtained. 
The  theory  of  dissociation  as  stated  by  Arrhenius  is  still  useful. 
Other  new  ideas  have  added  to  the  accuracy  of  our  concept  of  dis- 
sociation. For  example,  we  know  that  molecules  of  compounds  such 
as  carbon  dioxide  do  not  dissociate  in  solution  and  are  nonelectro- 
lytes  because  they  are  nonpolar.  However,  these  new  ideas  are  of 
value  chiefly  in  dealing  with  phenomena  which  are  beyond  the  scope 
of  this  book.  For  this  reason,  we  shall  follow  the  theory  of  ioniza- 
tion as  Arrhenius  originally  presented  it. 

What  happens  during  the  electrolysis  of  water?  So  far  we  have 
considered  pure  water  a  nonelectrolyte.  Roughly  speaking,  this  is 
true,  since  in  about  1  billion  grams  of  water  only  about  1  gram  of 
free  hydrogen  ions  is  present.  Yet  this  very  slight  dissociation  of  wa- 
ter [H,O  — »  H+  +  (OH)  -]  is  important.  Water  is  such  a  poor  con- 
ductor that  acid  must  be  added  to  conduct  the  current.  Nevertheless, 
the  hydrogen  and  oxygen  that  are  products  of  electrolysis  come  from 
the  water.  How? 

The  positive  hydrogen  ions  travel  to  the  negative  cathode.  The 
cathode,  connected  to  a  battery,  is  supplied  with  electrons.  When  the 
hydrogen  ion  reaches  the  cathode,  an  electron,  e,  from  the  cathode 
is  given  up  to  the  hydrogen  ion,  which  changes  to  a  neutral  hydro- 
gen atom.  This  atom  immediately  joins  another  hydrogen  atom  and 
is  liberated  as  a  molecule  of  gas,  thus: 

H+  +  c  -»  H°;     H°  +  H°  -»  H2  T 



The  negative  sulfate  ions,  and  hydroxyl  ions  from  the  water,  travel 
to  the  positive  anode.  The  hydroxyl  ion  gives  up  its  electron  more 
readily  than  does  the  sulfate  ion.  Having  lost  its  electron,  it  becomes 
the  hydroxyl  radical,  which  breaks  down,  forming  water  and  oxygen. 
The  sulfate  ions  remain  in  solution.  The  concentration  of  the  acid 
remains  unchanged. 

(OH)--€-»(OH);     4(OH)->2H20  +  02! 

What  is  hydrolysis?  We  should  expect  neutral  salts  to  show 
neither  acid  nor  basic  properties,  since  they  contain  neither  hydrogen 
nor  the  hydroxyl  radical  which  might  form  ions.  Yet  when  we  add 
blue  litmus  to  a  solution  of  copper  sulfate,  the  litmus  turns  pink,  in- 
dicating an  acid.  Where  are  the  free  hydrogen  ions  to  account  for 
this  behavior?  The  answer  lies  in  the  fact  that  water  is  slightly  dis- 

Cu++  +  (SO4)~  +  2HOH  «=»  2H+  +  (SO4)~  +  Cu(OH)2 

Since  the  copper  hydroxide  formed  is  only  very  slightly  dissoci- 
ated, there  will  be  some  union  of  Cu++  and  (OH)  -,  thus  forming  this 
very  weak  base,  and  liberating  an  excess  of  hydrogen  ions  that  form 
the  strong  acid,  sulfuric  acid.  Therefore,  the  solution* is  slightly  acid. 
This  reaction  is  the  reverse  of  the  neutralization  of  sulfuric  acid  by 
the  base  Cu  (OH) ,.  Hydrolysis  of  a  salt  is  the  action  between  the  ions 
of  water  and  the  ions  of  a  dissolved  salt,  forming  an  acid  and  a  base. 

Hydrolysis  also  explains  the  basic  nature  of  Na2CO3,  thus: 

2Na+  +  (C0*)~  +  2HOH  <=±  2Na+  +  2(OH)~  +  H2CO8 

Since  the  H2CO3  is  a  weak  acid,  there  will  be  some  union  of  H+ 
and  (CO8)  — ,  thus  forming  the  slightly  dissociated  H2CO3,  and  lib- 
erating an  excess  of  OH  ions  which  form  the  strong  base,  NaOH. 

A  solution  of  sodium  chloride  and  water  is  perfectly  neutral,  since 
it  is  made  from  a  strong  acid  and  a  strong  base.  A  solution  of  a  salt 
made  from  a  strong  base  and  a  weak  acid,  or  from  a  weak  base  and 
a  strong  acid,  does  not  show  a  neutral  reaction. 


Complete  the  following  equations,  and  tell  whether  the 
solution  in  each  case  will  be  acidic,  basic,  or  neutral  toward 


a)  K2CO3  +  2HOH -»  d)  Na2SO4  +  2HOH  -» 

b)  ZnSO4  +  2HOH-»  e)  KNO3  +  HOH  -> 

c)  A1C13  +  3HOH  ->  f)  K2SO4  +  2HOH  -» 

What  are  reversible  reactions?  Most  reactions  we  deal  with  are 
reactions  that  take  place  in  water  solutions  and,  hence,  are  reactions 
between  ions.  The  presence  of  free  ions  facilitates  chemical  changes. 
If  we  dissolve  sodium  chloride,  NaCl,  and  potassium  nitrate,  KNO3, 
in  water,  these  salts  immediately  dissociate,  forming  free  Na+,  Cl~, 
K+,  and  (NO3)  ~.  These  swiftly  moving  ions  constantly  meet  and 
form  molecules  of  KNO3  and  NaCl.  The  entire  reaction,  therefore, 
is  said  to  be  reversible,  that  is,  it  goes  in  both  directions. 

KNO3  +  NaCl  *±  KC1  +  NaNO3 

A  reversible  reaction  always  reaches  a  point  at  which  change  is  no 
longer  apparent.  In  other  words,  the  reaction  has  reached  a  point  of 
balance  or  equilibrium.  This  does  not  mean  that  nothing  is  happen- 
ing. On  the  contrary,  the  equilibrium  is  dynamic,  or  moving,  for  the 
substances  are  breaking  up  as  rapidly  as  they  are  being  formed. 

Reversible  reactions  and  equilibrium.  The  dissociation  of  an  elec- 
trolyte is  a  reversible  reaction.  Thus,  when  acetic  acid,  HC2H3O2,  is 
dissolved  in  water,  free  hydrogen  ions  and  free  acetate  ions  are 
formed.  These  ions  meet  and  form  acetic  acid. 

HC2H302  *=±  H+  +  (CzHsOz)- 

Finally  a  time  is  reached  when  the  rate  of  change  from  free  ions 
to  HC2H3O2  will  just  equal  the  rate  of  change  from  HC2H3O2  to 
free  ions.  This  will  happen,  in  the  case  of  a  0.0 IN  acid  solution, 
when  98.7  percent  of  the  HC2H3O2  is  in  the  form  of  free  ions  and  1.3 
percent  is  in  the  form  of  HC2H3O2.  When  this  condition  is  reached,  it 
would  appear  that  the  HC2H3O2  is  no  longer  dissociating  and  that  a 
condition  of  stable  equilibrium  has  been  reached.  In  reality  the  equi- 
librium is  not  stable  but  dynamic,  or  changing.  Changes  go  on  even 
after  the  98.7  :  1.3  ratio  is  attained,  but  while  new  HC2H3O2  is  being 
formed,  more  is  being  ionized,  keeping  the  same  ratio. 

When  do  reactions  go  to  completion?  Substances  that  do  not 
dissociate  cannot  react  reaoily  with  other  substances  that  do.  If  even 
one  of  the  products  of  a  reaction  is  unable  to  dissociate  to  any  great 
extent,  the  backward  action  cannot  take  place.  The  reaction  is  then 
said  to  go  to  completion.  Advantage  is  taken  of  this  to  secure  reac- 
tions that  complete  themselves. 

Fig.  52.  Action  of  silver  nitrate  and  dilute  hydrochloric  acid. 
Does  this  reaction  go  to  completion?  Why?  Under  what 
conditions  do  reactions  go  to  completion? 

Insoluble  substances,  called  precipitates,  gases  liberated  under  nor- 
mal temperature  conditions,  and  pure  water  are  practically  incapable 
of  dissociation.  Hence  a  reaction  will  go  to  completion  whenever  one 
of  the  products  formed  is  (1)  a  precipitate,  (2)  a  gas,  or  (2)  water. 
Examples  of  such  reactions  follow: 

1)  Formation  of  a  precipitate.  Chemists  use  as  tests  reactions  in 
which  precipitates  are  formed.  Thus,  in  testing  for  a  chloride,  silver 
nitrate,  AgNO;{,  is  added.  The  formation  of  insoluble  silver  chloride, 
AgCl,  prevents  the  reaction  from  reversing  itself. 

Ag+  +  (N03)-  +  Na+ 

AgCl  [  +  Na+  +  (NO8)- 

2)  Formation  of  a  gas.  IrTthe  laboratory  preparation  of  hydrogen 
chloride,  a  mixture  of  sodium  chloride  and  sulfuric  acid  is  heated, 
and  the  hydrogen  chloride  gas  that  is  liberated  leaves  the  field  of  ac- 
tion. Hence  the  reaction  does  not  reverse  itself. 


l-  +  H+  +  (HS04)  ~  -»  HC1 


3)  Formation  of  undissociated  water.  During  neutralization,  wa- 
ter is  always  one  of  the  products  formed.  Therefore  the  neutralizing 
reaction  is  complete,  since  water,  which  is  practically  undissociated, 
may  be  considered  as  having  left  the  field  of  chemical  action. 

Na+  +  (OH)-  +  H+  -f  Cl- -» Na+  -f  Cl~  +  HOH 
This  is  strictly  true  only  when  the  salt  formed  is  not  hydrolyzed. 


1.  Complete  and  balance  the  following  equations: 

a)  KNO3  +  NaCl  -»  c)  NaOH  +  PbCl2  -> 

b)  BaCl2  -f  K2SO4  -»  d)  KOH  4-  H2SO4  -» 



e)  Na2S03  +  H2S04  ->  SO,  +  H2O  + 

f)  FeS  4-  HC1  -> 

g)  CaCO3  -f  HN03  ->  CO2  +  H2O  + 
h)  CaCl2  +  NaN03  -> 

i)  Na2SiO3  -f  Ca(OH)2  -> 

2.  Examine  the  foregoing  equations  and,  in  each  case,  see 
whether  any  of  the  products  formed  are  gases  or  precipitates. 
Mark  them  with  the  appropriate  arrows,  j  or  j.  Remember 
that  S02,  H2S,  and  CO,  are  gases. 

3.  Finally,  with  the  aid  of  the  following  table  of  solubilities, 
tell  whether  each  of  the  reactions  goes  to  completion  or  not, 
stating  your  reason  in  each  case. 

TABLE  10. 

All  nitrates,  chlorates,  and  acetates  are  soluble  in  water.  SOLUBILITY  OF 

All  chlorides,  bromides,  and  iodides  are  soluble,  except  those  of  Ag,  Pb,  and  Hg.  JULUBILI  T  ur 

All  sul/ates  are  soluble,  except  PbS04,  BaS04,  CaSO,,  HgSO,,  and  Ag,S04.  COMMON 

All  hydroxides  are  insoluble,  except  those  of  Na,  K,  NH,,  Ca,  and  Ba.  COMPOUNDS 

All  sulfides  are  insoluble,  except  those  of  Na,  K,  NH4,  Ca,  and  Ba. 
All  oxides  are  insoluble,  except  those  of  Na,  K,  and  Ca. 
All  carbonates,  silicates,  and  phosphates  are  insoluble,  except  those  of  Na,  K,  and  NH.,. 

The  law  of  mass  action  and  equilibrium,  The  quantity  of  a  sub- 
stance in  a  unit  volume  of  solution  is  a  measure  of  the  concentration 
of  a  solution.  As  early  as  1803,  the  French  chemist  Claude  Berthollet 
(who  first  made  use  of  the  bleaching  action  of  chlorine)  noticed  that 
the  direction  of  a  chemical  reaction  is  dependent  upon  the  concentra- 
tions of  the  substances  involved  in  the  reaction. 

He  noticed  that,  in  general,  the  greater  the  concentration,  the 
greater  the  speed  of  the  reaction.  For  example,  a  match  burns  quietly 
in  ordinary  air  which  contains  about  21  percent  oxygen.  In  pure  oxy- 
gen, however,  the  match  burns  much  more  quickly,  since  the  con- 
centration of  the  oxygen  (one  of  the  reacting  substances)  has  been 
increased  almost  fivefold.  A  greater  concentration  of  oxygen  means 
more  molecules  of  oxygen  per  unit  volume  of  gas,  and,  therefore,  a 
greater  possibility  for  oxygen  molecules  to  come  in  contact  with 
molecules  of  carbon  and  carbon  compounds.  This  causes  an  increase 
in  the  speed  of  the  chemical  reaction  involved  in  burning. 

Sixty-four  years  later,  Guldberg  and  Waage,  professors  of  mathe- 
matics and  chemistry  at  the  University  of  Oslo,  Norway,  demon- 
strated that  the  speed  of  a  chemical  reaction  is  directly  proportional 
to  the  concentrations  of  the  reacting  substances.  They  also  made  an 
interesting  discovery  concerning  the  point  of  dynamic  equilibrium 



of  a  chemical  reaction  —  that  is,  the  poiixf  at  which  the  reaction 
in  one  direction  just  balances  the  reaction  in  the  opposite  direction. 
They  found  that  a  chemical  reaction  which  is  normally  reversible 
can  be  forced  to  go  in  one  direction  with  small  reversal.  This  is  ac- 
complished by  manipulating  the  concentrations  of  the  reacting  sub- 

Guldberg  and  Waage  expressed  this  phenomenon  in  the  form  of 
the  law  of  mass  action,  which  implies  that  a  change  in  the  quantity 
of  the  reacting  substances  results  in  a  change  in  the  equilibrium  point 
of  the  reaction.  In  the  manufacture  of  chemicals,  the  direction  of 
a  reaction  is  so  controlled  that  large  yields  are  produced. 

How  the  addition  of  a  common  ion  forces  a  reaction  to  go  to 
completion.  In  the  light  of  the  law  of  mass  action,  let  us  consider 
a  saturated  solution  of  sodium  chloride.  We  may  express  the  reac- 
tion that  is  taking  place  as: 

NaCl  <=>  Na+  +  Cl~ 

In  such  a  reaction,  the  product  of  the  concentrations  of  the  free  so- 
dium ions  and  the  free  chloride  ions  is  a  constant.  If,  by  any  means, 
we  increase  the  number  of  chloride  ions,  the  number  of  sodium  ions 
must  decrease.  The  number  of  sodium  ions  can  decrease  only  if  some 
of  the  sodium  chloride  comes  out  of  solution. 

To  increase  the  chloride  ions,  we  add  to  the  solution  a  compound 
of  chlorine  that  dissociates  to  a  high  degree.  Hydrochloric  acid  is 
such  a  compound.  Therefore,  if  hydrochloric  acid  is  added  to  a 
saturated  solution  of  sodium  chloride,  some  of  the  sodium  chloride 
will  be  precipitated.  We  call  this  shifting  of  the  equilibrium  point 
the  common-ion  effect.  In  this  case,  the  common-ion  effect  is  caused 
by  the  addition  of  the  chloride  ion,  which  is  common  to  HCi  and 

Fig.   53.   Apparatus   used  to 
show  the  common-ion  effect. 

NaCl  +  H2SO4 

J^  saturated 
NaCl  solution 

precipitated  out 

Svante  Arrhenius  (right)  and  his  close  friend,  Wilhelm  Ostwald.  Each  was  awarded 
a  Nobel  prize  for  chemistry,  Arrhenius  in  1903,  Ostwald  in  1909. 

Almost  insoluble  salts,  such  as  silver  chloride,  AgCl,  may  be  com- 
pletely precipitated  by  adding  a  common  ion.  Thus,  the  addition  of 
an  excess  of  NaCl  increases  the  concentration  of  the  few  dissociated 
chloride  ions  from  the  silver  chloride,  and  causes  some  more  AgCl 
to  precipitate. 

A  weak  acid  may  be  weakened  by  adding  a  salt  of  the  weak  acid. 
Thus  the  addition  of  sodium  citrate  to  citric  acid  weakens  that  acid, 
because  the  addition  of  the  common  ion  (the  citrate  ion)  forces 
more  of  the  citric  acid  to  the  undissociated  form. 


Jaffe,  Bernard.  Crucibles:  The  Story  of  Chemistry,  pp.  219- 
211.  Simon  and  Schuster,  New  York,  1948.  The  story  oi  the  life 
and  work  of  Arrhenius. 

Samrnis,  Constance  S.  "How  Annabella  Learned  the  Facts 
about  pH."  Journal  of  Chemical  Education,  Oct.,  1942,  pp.  490- 
494.  A  delightful,  cleverly  illustrated  treatment  of  a  very  im- 
portant topic. 


1.  Science  must  always  answer  at   least  two  questions  — 
how?  and  why? 

2.  Two  great  theories  underlie  much  of  chemistry:   (1)  the 
electron  theory  and   (2)  the  theory  of  dissociation. 



3.  The  theory  of  ionization  was  advanced  to  explain  elec- 
trolysis, nonelectrolytes,  strong  and  weak  acids  and  bases,  and 
other  puzzling  facts. 

4.  An  ion  is  an  atom  or  radical  that  carries  an  electric 

5.  When  an  acid,  base,  or  salt  goes  into  solution,  it  dis- 
sociates, partially  forming  free  ions.  When  a  current  is  passed 
through  such  a  solution,  the  ions  are  attracted  to  the  electrodes 
having  a  charge  opposite  their  own. 

6.  Free  ions  are  formed  in  greatest  numbers  in  water  so- 
lutions of  electrolytes.  Free  ions  occur  also  in  molten  sub- 
stances, in  heated  gases,  and  in  the  air  surrounding  radio- 
active substances.  Atoms  of  glowing  metals  throw  off  electrons; 
if  the  electrons  come  in  contact  with  a  gas,  as  in  a  radio  tube, 
the  molecules  of  the  gas  become  ionized. 

7.  In  terms  of  ionization,  an  acid  is  a  substance  that  pro- 
duces free  hydrogen  ions;  a  base  is  a  substance  that  produces 
free  hydroxyl  ions. 

8.  A  strong  acid  is  one  that  dissociates  easily,  forming  large 
numbers  of  hydrogen  ions.  A  weak  acid  forms  only  a  small 
number  of  hydrogen  ions.  Strong  bases  form  a  large  number  of 
hydroxyl  ions;  weak  bases  form  few.  A  compound,  such  as 
water,  whose  molecule  acts  like  a  tiny  magnet,  is  called  a 
dipole.  The  molecules  of  nonpolar  compounds,  such  as  CO2, 
do  not  dissociate  in  solution,  are  nonelectrolytes  and  noncon- 
ductors of  electricity. 

9.  Neutralization  is  the  union  of  the  hydrogen  ions  of  an 
acid  with  the  hydroxyl  ions  of  a  base. 

10.  The  acidity  or  alkalinity  of  a  solution  is  measured  in  pH 
values.  Water,  with  a  pH  of  seven,  represents  neutrality.  A  sub- 
stance with  a  pn  greater  than  seven  is  basic;  one  with  a  pn  less 
than  seven  is  acidic. 

11.  Substances   dissociate   to  different   degrees,   depending 
on  the  (1)  solute,   (2)   solvent,  (3)   concentration  of  the  solu- 
tion, and  (4)  temperature. 

12.  All  water  solutions  of  acids,  bases,  and  salts  are  elec- 
trolytes; that  is,  they  dissociate  and  conduct  an  electric  current. 

13.  Electrolytes   raise   the   boiling   points   and   lower    the 
freezing  points  of  solutions  to  a  greater  degree  than  non- 
electrolytes  do,  because  they  dissociate,  forming  larger  num- 
bers of  ions. 

14.  The  electron  theory  upholds  the  theory  of  dissociation. 
The  charge  on  an  ion  is  the  same  as  the  valence  of  the  element. 

15.  Hydrolysis  of  a  salt  is  the  action  between  the  ions  of 
water  and  the  ions  of  a  dissolved  salt,  forming  an  acid  and  a 
base.  It  is  the  opposite  of  neutralization.  Salts  formed  by  the 


reaction  of  a  strong  acid  and  a  weak  base  are  acidic;  by  the 
reaction  of  a  weak  acid  and  a  strong  base  are  basic. 

16.  A  reversible  reaction  is  one  that  will  go  in  two  direc- 
tions depending  upon  the  conditions  of  the  reaction.  Since 
most  chemical  reactions  are  reactions  between  free  ions,  if  the 
ions  are  neutralized  or  removed,  the  reaction  goes  to  com- 

17.  Reactions  go  to  completion  when  one  of  the  products 
formed  is  (1)  a  precipitate,  (2)  a  gas,  or  (3)  water. 

18.  The  common-ion  effect  is  caused  by  adding  an  ion  iden- 
tical with  one  of  the  ions  of  a  compound  in  solution.  It  results 
in  a  partial  precipitation  of  that  compound. 


Group  A 

1 .  To  what  questions  did  Arrhenius  seek  an  answer? 

2.  (a)  What  is  an  electrolyte?    (b)  What  three  classes  of 
compounds  are  electrolytes?   (c)  Give  three  examples  of  non- 

3.  How  could  you  find  out  whether  a  solution  contained 
an  electrolyte? 

4.  Explain  in  terms  of  the  dissociation  theory  what  hap- 
pens when  NaCl  is  dissolved  in  H^O. 

5.  Name  two  ways  in  which  ions  differ  from  atoms. 

6.  In  terms  of  the  theory  of  dissociation  define  (a)  an  acid, 
(b)  a  base. 

7.  (a)  On  what  do  the  properties  common  to  bases  de- 
pend?  (b)  to  acids? 

8.  Why  do  we   always  use  dilute  acid  in  preparing  hy- 
drogen  by  displacement  of  hydrogen   from   that   acid   by   a 

9.  Why  will  a  thoroughly  insoluble  hydroxide  not  turn 
pink  litmus  blue? 

10.  Explain  neutralization  in  terms  of  ions. 

1 1 .  Define  strong  acid  and  strong  base  in  terms  of  the  dis- 
sociation theory. 

12.  What  are  (a)  three  common  strong  acids,  (b)  two  com- 
mon weak  acids,    (c)  three  common  weak  bases,  and    (d)  two- 
common  strong  bases? 


13.  A  solution  has  a  pn  value  of  six.  What  does  this  mean? 

14.  Normal  blood  has  a  pn  of  about  7.3.  Is  it  acid  or  alka- 

15.  How  does  the  degree  of  dissociation  of  NH4OH  com- 
pare with  that  of  NaOH  in  equivalent  solutions? 

16.  (a)  What  happens  to  a  metal  atom  when  it  becomes  an 
ion?    (b)  Explain  the  dissociation  of  NaCl  in  terms  of  elec- 

17.  Explain  the  difference  in  physical  and  chemical  prop- 
erties of  the  potassium  atom  and  the  potassium  ion. 

18.  By  means  of  diagrams,  show  the  difference  between  the 
chloride  ion  and  the  chlorine  atom. 

19.  Change  the  following  equations  into  ionic  equations. 
Consult  the  table  of  solubilities  on  page  243. 

a)  AgNO3  +  KC1  -*  AgCl  +  KNO8 

b)  2NaCl  +  H2SO4  -»  2HC1  +  Na2SO4 

c)  BaCl2  +  H2S04  -»  BaSO4  +  2HC1 

20.  (a)  What  is  a  concentrated  acid?   (b)  Would  a  concen- 
trated acid  necessarily  be  a  strong  acid?  Explain. 

21.  Insoluble  bases  are  very  weak  bases.  Explain. 

22.  Complete  and  balance  the  following  equations. 

a)  Na2C03  +  CaCl2  ->  f)  Pb(NO3)2  +  NaCl  -» 

b)  Zn  +  H2S04-»  g)  BaCl2  +  K2SO4  -> 

c)  AgNO8  +  NaBr  ->  h)  NaNO3  -f  CuCl2  -> 

d)  NaCl  +  H2SO4  ->  i)  Pb(NO3)2  +  H2S  -> 

e)  NaCl  +  Zn(NO3)2  ->  j)  NaOH  +  NH4C1  -» 

23.  Which  of  the  reactions  in  exercise  22  are  reversible  and 
which  go  to  completion? 

24.  (a)  Which  of  the  following  salts  are  neutral  in  solution? 
(b)  Which  give  an  acid  reaction?   (c)  basic  reaction?   (d)  Ex- 

•  (1)  NaCl         1(3)  A1C13     l/(5)Na3BO8       (7)  ZnSO4 
Na2CO3       (4)  KNO3       (6)  K2SO4        J£)  Na2SO8 

25.  What  is  the  percentage  composition  of  sodium  acetate, 



26.  Explain  the  electrolysis  of  H2O  in  terms  of  ions  and 

Group  B 

27.  Why  does  cold  dilute  H2SO4  attack  Zn,  although  cold 
concentrated  H2SO4  does  not? 

28.  Dry  cobalt  chloride  (CoCl2)  is  blue  in  color.  A  solution 
of  this  salt  in  water  is  pink.  Explain  why  this  color  change 

29.  In  what  three  respects  has  the  original  theory  of  disso- 
ciation been  modified? 

30.  Concentrated  H2SO4  is  a  poorer  conductor  of  electricity 
than  dilute  H2SO4.  What  does  this  indicate  with  reference  to 

31.  Explain  how  the  study  of  the  freezing  points  of  solutions 
led  to  the  theory  of  dissociation. 

32.  (a)  State  the  law  of  mass  action,    (b)  Explain  how  it 
works,  using  a  suitable  reaction  as  an  example. 

33.  What  is  a  dipole?  Illustrate  your  answer. 

34.  What  weight  of  "Prestone,"  C2H4  (OH)  2,  should  you 
use  in  a  3|  gallon  auto  radiator  to  protect  it  against  5°F. 
weather?  One  mole  of  C2H4  (OH)  2  lowers  the  freezing  point 
of  one   liter  of  water  from  0°C.   to    -1.87°C.   Density  of 
C2H4(OH)2  is  1.13. 


1.  Write  a  report  or  organize  a  class  discussion  on  the  im- 
portance of  pn  measurements  to  a  soil  (agricultural)  chemist 
or  to  a  medical  research  man.  Consult  your  teacher  of  agri- 
culture or  your  family  doctor. 

2.  Construct  a  laboratory  setup  as  shown  on  page  239.  Plate 
some  copper  onto  a  piece  of  steel.  The  steel  object  should  be 
the  cathode  and  the  copper  strip  the  anode.  The  electrolyte 
should  be  a  copper  sulfate  solution. 

3.  Take  an  old  radio  tube  apart  and  demonstrate  to  the 
class  how  electrons  are  liberated  ifrom  the  heated  filament,  and 
how  the  flow  of  electrons  passes  through  the  grid  toward  the 
positive  plate.  Consult  your  teacher  of  physics  or  a  radio 
engineer  in  connection  with  this  project. 




The  men  of  experiment  are  like 
the  ant:  they  only  collect  and  use; 
the  reasoneis  resemble  spiders  who 
make  cobwebs  out  of  their  own  sub- 
stance. But  the  bee  takes  a  middle 
course;  it  gathers  its  material  from 
the  flowers  of  the  garden  but  trans- 
forms it  by  a  power  of  its  ozvn. 
Francis  Bacon  (1561-1626)  in  No- 

The  electron  structure  of  ammonia.  At  high  temperatures,  nitro- 
gen combines  with  hydrogen,  forming  ammonia,  NHV  Nitrogen  acts 
like  a  nonmetal  with  a  valence  of  three.  Since  its  outer  ring  is  short 
three  electrons,  it  combines  with  three  hydrogen  atoms,  iorming  a 
nonionic  compound.  ) 

The  three-dimensional,  or  cubic,  diagram  (Fig.  54)  is  another 
method  of  showing  the  arrangement  of  electrons  around  the  nu- 
cleus of  an  atom.  This  diagram  represents  the  formation  ot  ammonia. 
Note  that  the  electrons  of  the  three  hydrogen  atoms  are  shared  by 
the  nitrogen  atom  in  forming  the  ammonia  molecule.  Since  NH, 
bears  two  unshared  electrons,  it  can  accept,  or  combine  with  a  pro- 
ton (H+)  to  form  NH4+.  Hence,  it  is  a  base. 

How  ammonia  is  prepared  in  the  laboratory.  When  ammonium 
chloride  is  heated  with  calcium  hydroxide,  ammonia  gas  is  liber- 
ated and  may  be  collected  by  the  displacement  of  air  (Fig.  55) . 

2NH4C1  +  Ca(OH)2  ->  CaCl2  +  2NH3 1  +  2H2O 

Since  the  only  function  of  the  ammonium  chloride  is  to  furnish  the 
ammonium  group,  NH4,  almost  any  other  ammonium  salt  may  be 




Ammonia  NH3 

Fig.  54.  Formation  of  ammonia,  showing  electron  transfer* 

substituted  for  it.  Since  any  soluble  base  can  supply  the  OH  radical, 
we  can  use  any  soluble  base  instead  of  calcium  hydroxide.  The  gen- 
eral method  for  preparing  ammonia  is,  therefore,  by  the  reaction  of 
an  ammonium  salt  with  any  soluble  base. 

Physical  properties  of  ammonia.  Priestley  collected  ammonia  over 
mercury,  because  it  is  soluble  in  water.  At  room  temperature  1  vol- 
ume of  water  will  dissolve  more  than  700  volumes  of  this  gas,  NH.{. 

The  extreme  solubility  of  ammonia  may  be  shown  by  the  ammonia 
fountain.  This  setup  consists  of  a  flask  and  a  tube  (Fig.  56)  .  The 
flask  is  filled  with  dry  ammonia  gas,  and  inverted  over  water.  As 
soon  as  the  tube  enters  the  water,  some  ammonia  dissolves  in  it, 
reducing  the  pressure  inside  the  flask.  The  air  pressure  outside  then 
forces  the  water  up  into  the  tube,  and  as  it  issues  from  the  top  of 
the  tube  it  forms  a  fountain  of  ammonia  water.  Why? 

The  characteristic  pungent  odor  of  ammonia  was  known  to 
Priestley.  It  is  reported  that  as  he  bent  over  the  fireplace  where  he 
prepared  the  gas  by  the  method  we  use  today,  its  vapor  made  his 
eyes  fill  with  tears  and  drove  the  occupants  of  his  house  out  of  doors. 

Ammonia  gas  is  about  half  as  heavy  as  air.  It  may  be  readily  lique- 
fied, using  only  70  pounds  of  pressure  per  square  inch  at  ordinary 
temperatures.  The  colorless  liquid  NH3  is  kept  in  steel  cylinders  and 
shipped  in  tank  cars. 

Fig.  55.  Laboratory  preparation  of  ammonia. 
Why  is  the  generating  tube  tilted  downward? 

Fig.  56.  The  ammonia  fountain. 





Chemical  properties  of  ammonia.  Ammonia  reacts  with  water, 
forming  ammonium  and  hydroxyl  ions.  The  latter  ions  account  for 
the  basic  characteristics  shown  by  the  water  solution  of  ammonia, 
aqua  ammonia. 

NH3  +  H20  *±  (NH4)+  +  (OH)  - 

Ammonia  gas  does  not  burn  in  air,  but  it  does  burn  in  pure  oxy- 
gen with  a  pale  greenish  flame. 

4NH3  +  5O2  -»  6H2O  +  4NO 

Dry  ammonia  does  not  unite  with  dry  hydrogen  chloride,  but  the 
presence  of  a  trace  of  water  causes  the  two  to  combine,  forming  a 
white  cloud  that  settles  out  as  ammonium  chloride  powder. 

NH3  +  HC1  <=t  NH4C1 

Ammonia  unites  with  the  very  active  metals.  When  passed  over 
hot  magnesium,  for  example,  magnesium  nitride  and  hydrogen  are 

2NH3  +  3Mg  ->  3H2  +  Mg3N2 

Ammonia  is  a  very  useful  compound.  Liquefied  ammonia  gas,  or 
anhydrous  ammonia  as  it  is  called,  is  a  widely  used  compound.  In  a 
recent  normal  year,  more  than  2,000,000  tons  were  produced  in  this 
country  for  use  in  the  chemical  industries. 

More  than  75  percent  of  all  chemical  nitrogen  products  are  pro- 
duced from  ammonia.  Ammonia  is  the  chief  raw  material  from 
which  nitric  acid  and  nitrates,  or  salts  of  nitric  acid,  are  made.  Com- 
bined with  acids  to  form  ammonium  sulfate,  (NH4)2SO4,  ammo- 
nium nitrate,  NH4NO3,  and  monoammonium  phosphate,  NH4H3PO4, 
ammonia  is  one  of  the  chief  sources  of  nitrogen  in  fertilizers.  Urea, 
used  as  a  fertilizer,  is  also  made  from  ammonia.  The  manufacture 



In  recent  years,  there  has  been  an  increasing  use  of  anhydrous 
ammonia  as  a  primary  fertilizer.  Sometimes  the  ammonia  is  added 
to  irrigation  water  which  effectively  spreads  it  throughout  the  soil 
A  more  widespread  practice  is  to  inject  liquid  anhydrous  ammonia 
directly  into  the  soil  through  tubes  mounted  on  a  plow-like  ap- 
plicator. Upon  release,  the  ammonia,  which  has  been  kept  under 
pressure,  reverts  to  a  gas  and  unites  chemically  with  the  soil  par- 
ticles. Greatly  increased  crop  yields  have  resulted  from  treating 
the  soil  with  anhydrous  ammonia. 

Ammonia  is  widely  used  as  a  refrigerant.  Dry  ammonia  gas  is  used 
extensively  in  commercial  refrigeration  because  of  the  ease  with 
which  it  can  be  liquefied.  In  making  artificial  ice,  ammonia  gas  is 
placed  m  a  closed  system  of  pipes  and  coils.  By  means  of  a  pump, 
the  gas  is  compressed  until  it  changes  to  a  liquid;  the  heat  evolved 
in  the  process  is  removed  by  a  spray  of  cold  water.  Then,  as  cold 
hquid  ammonia,  it  is  passed  through  pipes  into  the  freezing  cham- 
ber, which  contains  brine  or  a  water  solution  of  calcium  chlo- 
ride,  C_jaCvd9. 

As  fast  as  the  liquid  ammonia  enters  the  pipes  in  the  chamber 
through  a  needle  valve,  it  expands  suddenly  and  vaporizes  as  a  result 
ot  the  reduced  pressure,  and  in  so  doing  absorbs  a  great  amount  of 
heat  from  the  brine.  So  cold  does  the  brine  become  that  the  pure 
water  m  the  tanks  that  are  immersed  in  the  brine  changes  to  blocks 
of  ice.  The  ammonia  gas  returns  to  the  pumping  chamber,  where  it 
is  again  compressed  for  reuse.  The  process  is  continuous,  and  the 
same  ammonia  gas  is  used  over  and  over  again. 

Ammonia  gas  also  helps  purify  water.  One  method  of  water  puri- 
fication of  increasingly  wide  use  employs  both  chlorine  and  ammo- 
nia.  This  method  is  known  as  the  chloramine  process  and  depends 

cold  water 

liquid  ammonia-*- 

1   -—  .  •=•-•••:•"  "•  .-.-.•-^"  *• 


Fig.  57.  Ammonia  refrigeration  equipment. 

on  the  fact  that  when  ammonia  and  chlorine  react,  chloramine, 
NH2C1,  is  formed.  One  of  the  reactions  is: 

2NH,  +  C12  ->  NH2C1  -f  NH4C1 

Chloramine  is  a  very  effective  killer  of  bacteria,  or  bactericide,  but 
is  less  active  chemically  than  chlorine  and  produces  less  of  the  typical 
taste  of  chlorinated  water.  The  process  is  somewhat  less  expensive 
than  chlorination,  and  is  especially  well  adapted  to  communities 
whose  water  supplies  have  musty  or  swampy  tastes. 

Preparation  and  properties  of  ammonium  hydroxide.  Care  must 
be  taken  to  distinguish  between  ammonia,  NH3,  and  the  ammo- 
nium ion,  NH4+.  In  the  laboratory  preparation  of  ammonia,  the  ni- 
trogen compound  first  formed  might  be  thought  of  as  ammonium 
hydroxide,  which  breaks  down  into  ammonia  and  water. 

2NH4C1  +  Ca(OH)2  ->  CaCl2  -f  [2(NH4)+  +  2 (OH)-] 

2NH3  +  2H20 

A  water  solution  of  NH3  is  often  called  ammonium  hydroxide  even 
though  the  compound  probably  does  not  exist.  The  solution  really 
consists  of  some  dissolved  NH3,  some  NH4  ions,  and  some  OH  ions.  A 
saturated  solution  of  ammonia  is  lighter  than  water  (sp.  gr.  0.88) , 
and  contains  about  36  percent  NH8  by  weight  at  room  temperatures. 
The  ammonia  may  be  expelled  by  boiling.  Household  ammonia  is 
a  water  solution  of  NH3  containing  about  6  percent  oleic  acid. 

A  water  solution  of  ammonia  is  a  weak  base,  since  NH3  reacts  with 
water  to  form  only  a  few  (one  percent)  hydroxyl  ions.  Because  it  is 
a  base,  it  dissolves  grease  and  hence  removes  dirt.  Since  this  ammo- 
nia water  gives  off  vapor,  or  volatilizes,  rapidly  and  completely  with- 
out leaving  a  solid,  it  is  useful  as  a  household  cleansing  agent. 




The  ammonium  radical  and  ammonium  salts.  The  ammonium 
radical,  because  of  its  positive  valence,  is  considered  metallic.  Its 
presence  may  be  detected  by  adding  the  salt  to  be  tested  to  a  base, 
such  as  sodium  hydroxide,  and  heating  the  mixture.  If  the  sus- 
pected substance  is  an  ammonium  compound,  ammonia  gas  will  be 
liberated,  and  can  be  identified  by  its  odor.  Moist  pink  litmus  paper 
held  in  the  gas  turns  blue.  Why  must  the  litmus  be  wet? 

(NH4)2S04  4  2NaOH  -+  Na2SO4  4-  2H2O  4-  2NH3  T 

One  of  the  most  common  of  the  ammonium  salts  is  ammonium 
chloride,  or  sal  ammoniac.  This  salt,  first  produced  in  Egypt,  was 
known  to  the  early  alchemists.  It  is  a  white,  crystalline  substance, 
readily  soluble  in  water.  It  is  decomposed  by  heat  into  two  gases, 
NH3  and  HC1,  which  reunite  on  cooling. 

,,  NH4C1  *=>  NH3  4-  HC1 

Sal  ammoniac  is  used  extensively  in  dry  batteries  as  an  electrolyte 
and  in  soldering  in  which  the  hot  iron  dissociates  the  salt.  The  hydro- 
gen chloride  arid  ammonia  liberated  remove  the  rust  that  covers  the 
surface  of  the  metal  to  be  soldered,  the  hydrogen  chloride  by  dissolv- 
ing the  rust,  and  the  ammonia  by  reducing  it. 

Fritz  Haber  makes  the  synthesis  of  ammonia  a  commercial  suc- 
cess. Minor  quantities  of  ammonia  are  still  made  by  an  old  process  — 
the  destructive  distillation  of  coal.  This  process,  in  which  ammonia 
is  produced  as  a  valuable  byproduct,  is  discussed  on  pages  862— 363. 
Synthetic  methods  have  superseded  the  method  of  preparing  am- 
monia from  coal.  The  most  successful  synthesis  of  ammonia  is  based 

A  compressor  used  in  the  manufacture  of  ammonia.  The  hydogen  is  made  from 
natural  gas  produced  in  nearby  oilfields. 

Shell  Chemical  Corporation 


on  the  process  first  worked  out  on  a  commercial  basis  by  Fritz  Haber 
(ha'ber)  in  1913  and  known  as  the  Haber  process. 

The  first  real  test  of  this  great  achievement  came  during  World 
War  I,  partly  as  a  result  of  the  desperate  need  of  the  German  gov- 
ernment for  nitrogen  compounds.  Haber's  process  made  agriculture 
in  blockaded  Germany  independent  of  Chile  saltpeter  and  also  gave 
the  German  military  machine  a  new  source  of  nitrates  for  high  ex- 
plosives. The  Haber  process  enabled  Germany  to  fight  hunger,  and 
stave  off  defeat  much  longer  than  the  Allies  expected.  Fritz  Haber 
was  later  forced  into  exile  by  the  Nazis  and  died  in  Switzerland. 

The  Haber  process.  This  process,  with  its  many  modifications,  is 
the  most  important  single  process  for  the  fixation  of  nitrogen,  that  is, 
the  combining  of  the  free  nitrogen  of  the  air  with  other  elements  to 
form  useful  compounds.  The  process  is  based  upon  what  appears  to 
be  a  very  simple  reaction,  the  union  of  hydrogen  and  nitrogen  gases. 

The  nitrogen  and  hydrogen  for  this  reaction  are  obtained  from 
air,  coke-oven  gas,  water  gas,  natural  gas,  and  some  petroleum  refin- 
ery gases.  Although  this  chemical  reaction  has  been  known  for  a 
long  time,  it  was  not  industrially  practicable  until  the  reversible 
reaction  could  be  controlled.  The  laws  of  chemical  equilibrium  had 
to  be  used,  so  that  some  of  the  ammonia  gas  formed  would  not  be 
immediately  decomposed  into  its  elements. 

What  factors  can  be  used  to  control  the  point  of  equilibrium? 
Most  chemical  reactions  are  reversible.  In  the  synthesis  of  ammonia, 
the  first  equilibrium  ratio  was  about  two  percent  ammonia  to  98  per- 
cent of  a  mixture  of  nitrogen  and  hydrogen.  In  other  words,  most  of 
the  ammonia  formed  during  the  union  of  nitrogen  and  hydrogen  de- 
composed into  its  constituent  gases. 

The  point  of  equilibrium,  however,  can  be  controlled  to  some 
extent.  The  factors  that  help  to  control  it  are  (1)  temperature, 
(2)  pressure,  and  (3)  concentration  of  the  substances  involved  in 
the  reaction.  Catalytic  agents  increase  the  speed  of  a  reaction  and, 
hence,  enable  the  point  of  equilibrium  to  be  reached  more  rapidly. 
It  is  not  probable  that  catalysts  alter  the  point  of  equilibrium. 

Thus,  in  the  case  of  the  preparation  of  ammonia  in  America  by  a 
modified  Haber  process,  the  reaction 

3H2  +  N2  -*  2NH3  T  +  24,000  calories 

has  been  forced  to  go  to  the  right,  producing  as  much  as  30  percent 
NH3,  instead  of  only  two  percent  of  the  theoretical  yield. 



\itrogen  Division,  Allied  Chemicul  rind  Dye  Corporation 

Aerial  view  of  a  large  plant  for  the  synthesis  of  ammonia  and  the  manufacture 
of  fertilizer. 

The  conditions  that  made  the  American  process  a  success  were 
(1)  the  use  of  a  specially  prepared  iron  oxide  as  catalyst,  (2)  a  tem- 
perature of  about  475°C.,  (3)  a  reaction  pressure  of  about  300  atmos- 
pheres (atrn.) ,  that  is,  300  times  atmospheric  pressure,  and  (4)  a 
rapid  removal  of  the  NH3  formed.  Since  heat  is  evolved  during  the 
synthesis  of  ammonia,  the  higher  the  temperature,  the  less  the  yield, 
and  consequently  too  high  a  temperature  is  avoided.  Since  there  is 
a  diminution  of  volume  (three  volumes  of  hydrogen  unite  with  one 
volume  of  nitrogen,  forming  only  two  volumes  of  ammonia) ,  the 
yield  is  increased  by  an  increase  in  pressure.  The  manufacturing 
conditions  represent  the  most  effective  compromise  between  largest 
yield,  shortest  time,  and  most  profitable  rate. 

Another  synthetic  process  for  making  ammonia.  In  1916  the 
United  States  built  a  large  nitrogen-fixation  plant  at  Muscle  Shoals, 
Alabama,  where  ammonia  was  to  be  prepared  by  the  cyanamide  proc- 
ess. This  project  was  later  abandoned,  although  the  cyanamide 
process  at  Niagara  Falls,  Canada,  has  been  very  successful  as  a  source 
of  calcium  cyanamide.  The  chief  chemical  changes  that  take  place 
in  the  cyanamide  process  are: 

1)  The  formation  of  calcium  carbide  by  heating  lime,  CaO,  and 
coke,  C,  in  an  electric  furnace. 

CaO  +  3C  -»  CO  -j-  CaC2  (calcium  carbide) 


2)  The  union  of  calcium  carbide  with  free  nitrogen,  forming 
calcium  cyanamide. 

CaC2  +  N2  — >•  C  +  CaCN2  (calcium  cyanamide) 

3)  The  addition  of  steam  to  cyanamide,  forming  ammonia. 

CaCN2  +  3H2O  -»  CaCO8  +  2NH3 f 

In  dry  form,  crude,  powdered  calcium  cyanamide  containing  about 
60  percent  CaCN2  is  sold  as  a  fertilizer  under  the  name  of  "Cyana- 
mid."  All  of  its  nitrogen  is  available  as  a  plant  food. 


Berl,  E.  "Fritz  Haber."  Journal  of  Chemical  Education,  May, 
1937,  pp.  203-207.  A  short  biography. 

Jaffe,  Bernard.  Men  of  Science  in  America,  pp.  307-330. 
Simon  &  Schuster,  New  York,  1944.  The  development  of  our 
ideas  regarding  reversible  reactions  is  tied  up  with  the  contri- 
bution of  America's  greatest  theoretical  scientist,  J.  Willard 
Gibbs.  His  life  and  work  are  described  here. 

Ross,  William  H.;  Adams,  J.  Richard;  Yee,  J.  Y.;  and  Whit- 
taker,  Colin  W.  "Preparation  of  NH4NO3  for  Fertilizer  Use." 
Industrial  and  Engineering  Chemistry,  Dec.,  1944,  pp.  1088- 

Slosson,  Edwin  E.,  Creative  Chemistry,  pp.  14-36.  D.  Apple- 
ton-Century  Co.,  New  York,  1920.  "Nitrogen,  Preserver  and 
Destroyer  ol  Lite." 


1.  When  a  substance  which  ionizes  goes  into  solution,  the 
change  that  takes  place  is  reversible.  The  substance  dissociates 
into  ions;  these  ions  unite,  re-forming  the  original  substance. 
Even  when  equilibrium  is  established,  this  reversible  reaction 
continues.  Change  is  no  longer  apparent,  however,  because 
the  rate  of  dissociation  is  the  same  as  the  rate  at  which  the  ions 
in  the  solution  reunite,  forming  the  original  substance.  The 
solution  is  in  a  state  of  dynamic  equilibrium. 

2.  The  point  of  equilibrium  ot  a  reversible  reaction  can  be 
controlled  to  a  certain  extent  by    (1)  temperature,    (2)  pres- 
sure, and   (3)  concentration  of  the  substances  involved.  Cata- 
lytic agents  increase  the  speed  of  a  reaction  and,  hence,  enable 
the  point  of  equilibrium  to  be  reached  more  rapidly. 



Group  A 

1.  Describe   the   laboratory   preparation   of  NH3.   Use   a 
labeled  diagram. 

2.  (a)  How  did  Priestley  first  collect  ammonia?    (b)  Why 
did  he  use  this  method? 

3.  (a)  What   is   the  general   method   of  preparing  NH3? 
(b)  Write  two  equations  illustrating  two  ways  of  preparing 
NH3  from  NH4C1. 

4.  Show,  by  a  simple  experiment,  the  extreme  solubility  of 
NH,  in  H2O. 

5.  Devise  a  simple  experiment  to  determine  whether  NH8 
is  lighter  or  heavier  than  air. 

6.  Does  NH3  burn?  Explain. 

7.  (a)  By  means  of  an  equation,  give  the  chief  chemical 
property  of  NH3.   (b)  What  type  of  reaction  is  this?   (c)  How 
can  you  make  the  reaction  go  to  the  right?  (d)  to  the  left? 

8.  Complete  and  balance  the  following  equations: 

a)  NH3  +  HC1  -»  c)   (NH4),S04  +  NaOH-> 

b)  NH3  +  Mg  ->  d)  NH3  +  H2SO4 


9.  What  are  five  uses  of  NH3? 

10.  What  substances  are  present  in  a  water  solution  of  NH3? 

11.  (a)  Is  a  solution  of  NH3  in  water  a  strong  base?  (b)  Ex- 
plain your  answer. 

12.  (a)  Write  an  ionic  equation  showing  the  neutralization 
of   HC1   by   NH4OH,    (b)  also   the   neutralization   of   H2SO4 
and  (c)  of  HNO3  by  the  same  base. 

13.  Does  dry  NH3  affect  litmus?  Explain. 

14.  (a)  In  what  group  of  elements  is  the  ammonium  radical 
placed?  (b)  Explain. 

15.  What  are  two  commercial  methods  of  preparing  NH3? 

16.  What  was  the  difficulty  that  confronted  manufacturers 
who  attempted  to  make  NH3  by  direct  synthesis  of  its  ele- 

17.  What  four  conditions  are  met  in  the  improved  Haber 


18.  (a)  Describe  the  cyanamide  process  for  making  NH8. 
(b)  What  are  the  three  equations? 

19.  Complete  and  balance  the  following  equations: 

a)  NH3  +  HNO3  ->  b)  NH3  +  H3PO4  -> 

20.  "Spirits  of  hartshorn"  was  the  name  applied  to  am- 
monia water  prepared  by  the  alchemists  by  heating  the  horns 
of  deer.  What  elements  must  have  been  present  in  deer  horns? 

21.  In  what  two  ways  do  liquid  ammonia  and  aqua  am- 
monia differ? 

22.  What  simple  test  will  distinguish  NH^Cl  from  Nad? 
Both  of  these  compounds  are  white,  soluble  salts. 

23.  Compare  the  laboratory  methods  of  collecting  NH3,  N2, 
HC1,  02,  C12,  and  H2. 

24.  Explain  the  action  of  liquid  NH3  on  dry  litmus  paper. 

25.  Why  is  ammonia  water  called  the  volatile  alkali? 

26.  Bottles  of  household   ammonia  were   formerly  closed 
with  rubber  rather  than  with  glass  or  cork.  Why? 

27.  Why  is  ammonia  water,  rather  than  lye,  used  in  re- 
moving grease  spots  from  clothing? 

28.  Compare  the  ease  with  which  NH3,  N2,  C12,  and  H2  are 

29.  Describe  the  preparation  of  artificial  ice  by  means  of 
NH3.   (b)  What  property  of  NH8  makes  it  useful  in  this  proc- 
ess? (c)  Of  what  use  is  the  CaCl0  or  brine  solution? 


30.  What  weight  of  NH4C1  is  necessary  to  make  340  tons 
of  dry  liquid  NH3? 

31.  Determine  the  percentage  of  nitrogen  in  monoammon- 
ium  phosphate. 

32.  (a)  What  is  chloramine?    (b)  How  is  it  prepared? 

Group  B 

33.  Bottles  of  ammonia  water  and  hydrochloric  acid  are 
placed  within  a  few  inches  of  each  other  and  their  stoppers  re- 
moved. White  fumes  are  seen.  Explain. 


34.  Explain  how  you  would  determine  the  strength  of  a  bot- 
tle of  household  ammonia  by  titration.  Give  details. 

35.  How  can  a  knowledge  of  the  laws  of  equilibrium  be 
used  in  making  the  preparation  of  NH8  by  the  Haber  process 
more  efficient  and  more  economical? 

36.  When  water  is  added  to  magnesium  nitride   (Mg3N2) , 
the  odor  of  NH3  is  detected.  Write  an  equation  to  explain  this. 

37.  CaCl2  unites  with  NH3,  forming  CaCl2  •  8NH3.  Explain 
whether  you  could  use  CaCl2  in  drying  ammonia  gas. 

38.  Account  for  the  odor  of  NH3  around  heaps  of  garbage 
and  manure. 


1.  Take  a  dry  cell  apart,  and  show  the  various  substances 
that  are  used  in  its  manufacture.  Prove  the  presence  of  NH4C1 
by  a  chemical  test.  What  is  the  purpose  of  the  NH4C1? 

2.  Write  a  two-  or  three-page  report  on  the  freedom  of  the 
man  of  science  in  a  democracy  compared  with  the  enslavement 
of  science  in  totalitarian  states.  Use  Haber  and,  perhaps,  Lang- 
muir  as  examples. 

3.  Set  aside  2  ml.  of  household  ammonia  in  an  open  test 
tube  for  two  weeks.  Test  for  the  presence  of  ammonium  hy- 
droxide with  pink  litmus  paper  at  the  end  of  the  time.  Explain 
the  results. 




.  .  .  For  nitrogen  plays  a  double  role 
in  human  economy.  It  appears  like 
Brahma  in  two  aspects,  Vishnu  the 
Preserver  and  Siva  the  Destroyer. 
E.  E.  Slosson,  1919 

The  revolution  brought  about  by  man-made  nitrogen  com- 
pounds. Haber's  successful  synthesis  of  ammonia  widened  man's 
control  over  nature  by  making  him  tess  dependent  for  his  raw  mate- 
rials on  supplies  present  in  limited  or  faraway  areas  of  the  earth's 
surface.  Synthetic  ammonia  was  soon  converted  into  nitric  acid. 
This,  in  turn,  gave  mankind  an  unlimited  supply,  not  only  of  ferti- 
lizers, but  also  of  high  explosives.  The  synthesis  of  nitrogen  com- 
pounds from  air  and  water  was  a  bloodless  revolution  whose  conse- 
quences touched  the  lives  of  half  the  people  of  the  world. 

Preparation  and  properties  of  nitric  oxide.  Not  only  does  nitrogen 
unite  with  hydrogen  at  high  temperatures,  but  it  combines  with  oxy- 
gen also  when  an  electric  spark  is  passed  through  a  mixture  of  the 
two  gases.  Cavendish  made  this  discovery  in  1770  when  he  passed  an 
electric  spark  through  mixtures  of  hydrogen  and  air  in  his  synthe- 
sis of  water.  Soon  afterward  Priestley  made  a  thorough  study  of  the 
compound  formed  by  the  direct  union  of  nitrogen  and  oxygen. 

N2  +  O2  — »  2NO  (nitric  oxide) 

Priestley  also  prepared  nitric  oxide  by  the  action  of  copper  on 
dilute  nitric  acid,  collecting  the  gas  by  the  displacement  of  water, 




as  shown  in  the  illustration  below.  This  is  the  laboratory  method 
used  today. 

3tu  +  8HNO3  ->  3Cu(NO3)2  4-  4H2O  +  2NO  T 

Nitric  oxide  is  a  poisonous,  colorless  gas,  very  slightly  soluble  in 
water,  and  about  as  heavy  as  air.  Chemically,  it  is  very  active.  When 
exposed  to  air  or  oxygen,  it  oxidizes  at  once  to  nitrogen  dioxide,  NO2. 
This  reaction  produces  much  heat;  that  is,  it  is  decidedly  exothermic. 

Nitrogen  dioxide.  The  equation  representing  the  oxidation  of 
nitric  oxide  to  nitrogen  dioxide  is: 

2NO  +  O2 


Nitrogen  dioxide  is  reddish  brown  in  color,  heavier  than  air,  very 
soluble  in  water,  and  easily  liquefied.  Its  fumes  are  irritating  and 

In  preparing  nitric  oxide  by  means  of  copper  and  nitric  acid,  the 
gas  first  seen  in  the  generating  bottle  is  brown  nitrogen  dioxide 
rather  than  the  colorless  nitric  oxide.  Actually  NO  is  formed  first, 
but  combines  with  the  oxygen  in  the  air  of  the  generating  bottle  to 
form  the  brown  NCX.  When  the  mixture  of  the  NO  and  NO2  passes 
through  the  delivery  tube  into  the  collecting  bottle,  the  NO2  dis- 
solves in  the  water.  Only  the  colorless  nitric  oxide  displaces  the 
water  in  the  bottle. 

As  its  temperature  is  lowered,  nitrogen  dioxide  gradually  changes 
into  nitrogen  tetroxide  gas,  N2O4.  The  reaction  2NO2  ^  N2O4  is  a 
reversible  one.  Above  140°C.,  the  reaction  goes  to  completion,  as 
shown  by  the  arrow  pointing  to  the  left.  At  room  temperature,  the 
gas  obtained  is  a  mixture  of  the  reddish  brown  NO2  and  the  color- 
less N2O4.  At  very  low  temperatures,  NO2  changes  completely  into 

The  arc  process  of  making  nitric  acid  from  air  and  water.  The 
union  of  nitrogen  with  oxygen  in  an  electric  arc,  or  by  the  action  of 
an  electric  spark,  was  used  as  the  basis  of  many  commercial  ventures 
attempting  to  produce  synthetic  nitric  acid.  In  the  beginning,  they 

nitric  oxide 


Fig.  58.  Laboratory  prepara- 
tion of  nitric  oxide.  What  other 
laboratory  setup  does  this  re- 

copper  wire 
inside  glass 


Fig.  59.  Laboratory  preparation  of  nitrogen  dioxide  by  the  arc  process.  How  can 
you  tell  whether  NO2  is  formed? 

all  failed  because  when  a  mixture  of  nitrogen  and  oxygen  was  thus 
treated,  the  yield  of  nitric  oxide  was  very  small.  Because  the  reac- 
tion is  reversible,  practically  all  the  nitric  oxide  formed  at  first  was 
decomposed  into  its  original  elements. 

N2  4-  O2  ?=±  2NO  T 

A  careful  application  of  the  laws  of  chemical  equilibrium  finally  re- 
sulted in  larger  yields  of  nitric  oxide.  This  gas  combined  with  the 
oxygen  of  the  air  to  form  nitrogen  dioxide  which  was  then  dissolved 
in  water  to  form  nitric  acid,  HNO,. 

2NO  +  02 

3N02  +  H20 


>  2HN03  4-  NO 

The  Ostwald  process,  another  commercial  method  of  making 
nitric  acid.  The  arc  process  because  of  its  higher  cost  has  been  super- 
seded by  the  Ostwald  process.  No  sooner  had  the  synthesis  of  ammo- 
nia been  successfully  carried  out  than  Wilhelm  Ostwald  (ost'valt) , 
a  chemist  who  helped  Arrhenius  establish  the  theory  of  ionization, 
showed  how  ammonia  gas  could  be  converted  into  nitric  acid  on  an 
industrial  scale. 

The  Ostwald  process  consists  of  oxidizing  ammonia  gas  in  the  pres- 
ence of  a  catalyst,  which  consists  of  a  specially-prepared  platinum  or 
platinum-rhodium  gauze  heated  to  a  red  heat.  The  two  oxides  of 
nitrogen  are  formed  during  the  process,  but  the  final  products  may 
be  represented  by  the  following  equation: 

NH3  +  2O2  ->  HN03  +  H2O 

Using  a  pressure  of  only  about  six  atmospheres,  about  95  percent 
of  the  ammonia  is  converted  into  a  water  solution  of  nitric  acid  of 
about  50  percent  concentration. 

Nitrogen  fixation.  The  change  of  free  nitrogen  into  ammonia  by 
the  Haber  process  and  the  change  of  free  nitrogen  from  the  air  into 
nitric  acid  by  the  arc  process  are  examples  of  nitrogen  fixation. 


gauze"  1  combustion 




Fig.  60.  Laboratory  preparation  of  NOj  by  the  Ostwald  process. 

Nitrogen  fixation  is  the  changing  of  free  nitrogen  from  the  air  into 
useful  compounds.  During  electric  storms,  some  nitric  acid  is  formed 
in  the  atmosphere  by  a  natural  arc  process. 

Nitrogen-fixing  bacteria  help  the  farmer.  Most  farm  crops  use 
part  of  the  valuable  salts  in  soil.  On  the  other  hand,  certain  crops 
such  as  peas,  beans,  and  alfalfa,  actually  enrich  the  soil  in  which 
they  are  grown.  Chemistry  explains  the  fertility  of  such  soil  in  the 
following  way:  Plants  require  nitrogen  in  the  form  of  nitrates. 
These  nitrates  are  soluble  in  water  and  can,  therefore,  be  absorbed 
by  the  root  hairs  of  plants  by  osmosis,  a  process  by  which  liquids  and 
gases  pass  through  semipermeable  tissues.  On  the  roots  of  plants  such 
as  peas,  beans,  and  alfalfa  are  little  nodules,  inside  which  live  bil; 
lions  of  nitrogen-fixing  bacteria  (Rhizobium)  .  These  bacteria  and 
several  others  have  the  ability  to  change  the  free  nitrogen,  present 
in  porous  soil,  into  soluble  nitrogen  compounds,  particularly  ni- 
trates, that  plants  use  in  building  living  tissues.  When  such  crops 
are  plowed  under,  this  "green  manure"  adds  nitrogen  to  the  soil. 

The  nitrogen  cycle  —  nitrogen  compounds  break  down  into  free 
nitrogen.  Various  other  bacteria  break  down  nitrogen  compounds  in 
the  soil  into  simpler  compounds  and  even  return  considerable  quan- 
tities of  free  nitrogen  to  the  air.  Such  bacteria,  called  denitrifying 
(de-ni'trMI-ing)  bacteria,  cause  the  loss  of  nitrogen  from  the  soil  and 
thus  complete  the  extremely  important  nitrogen  cycle.  These  bacteria 
are  responsible,  in  part,  for  the  rapid  decay  of  nitrogen-containing 

These  pea  roots  are  covered 
with  nodules  containing  nitro- 
gen-fixing bacteria. 


organic  wastes.  They  are  used  widely  in  the  treatment  of  sewage  (see 
pages  222-223)  .  This  series  of  changes  is  referred  to  as  the  nitrogen 

How  nitric  acid  is  prepared  in  the  laboratory.  The  laboratory 
preparation  of  nitric  acid  follows  the  general  method  for  preparing 
an  acid.  Sodium  nitrate,  mixed  with  concentrated  sulfuric  acid,  is 
heated  gently  in  a  glass  retort.  Nitric  acid  is  formed,  which,  having 
a  lower  boiling  point  than  sulfuric  acid,  evaporates.  It  is  then  con- 
densed into  a  colorless  liquid  by  cooling,  as  shown  in  the  illustration 
below.  The  equation  for  the  preparation  of  nitric  acid  is: 

NaNO3  +  H2SO4  ->  HNO3  +  NaHSO4  (sodium  hydrogen  sulfate) 

This  method  is  also  one  of  the  commercial  processes  used  today  in 
making  about  10  percent  of  the  nitric  acid  consumed  by  the  world's 
industries.  Synthetic  NaNO3  supplies  the  nitrate. 

Nitric  acid  was  known  to  the  alchemists  more  than  1000  years  ago. 
Geber  (ga'ber)  ,  an  Arabian  physician  and  alchemist,  prepared  it 
about  A.D.  800.  It  was  called  aqua  fortis,  meaning  strong  water. 

Physical  properties  of  nitric  acid.  When  pure,  nitric  acid  is  a  color- 
less liquid.  Its  water  solution,  containing  68  percent  nitric  acid  by 
weight  (that  is,  the  concentrated  nitric  acid  of  commerce)  ,  has  a 
specific  gravity  of  1.4,  and  boils  at  120°C.  The  concentrated  acid 
fumes  strongly.  The  yellowish  appearance  of  the  nitric  acid  prepared 
in  the  laboratory  is  caused  by  the  presence  of  nitrogen  dioxide, 
formed  by  the  partial  decomposition  of  the  nitric  acid  during  the 

Chemical  properties  of  nitric  acid.  Because  nitric  acid  mixes  with 
water  in  all  proportions  and  dissociates  almost  completely,  thus  pro- 
ducing large  quantities  of  hydrogen  ions,  it  is  a  strong  acid.  Nitric 
acid  is  unstable.  In  sunlight  or  when  heated,  it  decomposes  into 
water,  oxygen,  and  nitrogen  dioxide. 

4HN03  --»  2H20  +  4NO2  T  +  O2  T 
Fig.  61.  Laboratory  preparation  of  nitric  acid.  By  what  process  is  the  acid  collected? 

cold  water 

nitric  acid 


A  glowing  splint  inserted  in  the  vapors  of  boiling  concentrated 
nitric  acid  catches  fire,  thus  showing  that  oxygen  is  present  in  the 
vapors.  When  an  element,  such  as  nitrogen  in  HNO3,  is  in  a  very 
highly  oxidized  state,  that  is,  it  has  a  high  positive  charge,  the  com- 
pound is  a  strong  oxidizing  agent. 

The  action  of  nitric  acid  on  metals  illustrates  its  oxidizing  pow- 
ers. When  hydrochloric  acid  reacts  with  many  of  the  metals,  hydro- 
gen is  liberated  even  if  concentrated  acid  is  used.  On  the  other  hand, 
dilute  nitric  acid  acts  on  a  metal,  forming  water  instead  of  hydrogen. 
Nitric  oxide  gas  is  also  formed.  In  fact,  one  of  the  methods  used  to 
prepare  nitric  oxide  depends  on  this  action  of  nitric  acid. 

3Cu  +  8HNO3  ->  3Cu(NO3)2  +  4H2O  +  2NO  j 

When  concentrated  nitric  acid  reacts  with  a  metal,  nitrogen 
dioxide,  instead  of  nitric  oxide,  is  formed.  Brown  NO2  is  produced 
by  the  oxidation  of  NO  to  NO.. 

Cu  +  4HN03  -»  Cu(N03)2  +  2H2O  +  2NO2 1 

The  electron  theory  explains  the  oxidizing  power  of  nitric  acid  as 
follows:  oxidation  is  a  loss  of  electrons;  therefore,  a  substance  such 
as  chlorine,  which  borrows  electrons,  is  a  good  oxidizing  agent. 
HNO3  may  be  thought  of  as  containing  N+++++ ,  which  borrows  elec- 
trons from  the  copper  in  the  reaction  above,  changing  to  N++  (in 
NO) .  Thus  the  copper  is  said  to  be  oxidized  and  the  nitrogen  re- 

Nitric  acid  attacks  proteins,  forming  a  yellow  compound.  The 
yellow  coloration  produced  on  the  skin  by  nitric  acid  is  caused  by 
this  chemical  reaction  (see  also  Chapter  36) .  Nitric  acid  oxidizes 
both  cotton  and  wool,  destroying  most  of  the  strength  of  the  fibers. 

Aqua  regia,  the  acid  mixture  that  dissolves  gold  and  platinum. 
Nitric  acid,  when  mixed  with  hydrochloric  acid,  oxidizes  the  latter, 
liberating  atomic  chlorine.  This  mixture  of  nitric  and  hydrochloric 
acids  as  in  the  equation  below  is  called  aqua  regia,  or  royal  water. 

3HC1  +  HNO3  -»  2H2O  +  NO  +  3C1 

This  chlorine  reacts  with  gold,  forming  soluble  gold  chloride. 

Au  +  3C1  ->  AuCl3 

Both  gold  and  platinum  are  insoluble  in  any  one  of  the  three  com- 
mon strong  acids,  but  they  are  soluble  in  aqua  regia. 


FeSO4  +  nitrate  solution  •    -j^nar  i 

jjjji&r  concentrated 

brown  ring--"^^ H2SO4 

Fig.  62.  The  brown-ring  test  for  a  nitrate. 

How  do  we  test  for  the  nitrate  ion?  The  nitrates,  salts  of  nitric 
acid,  are  all  soluble  in  water,  are  decomposed  by  heat,  and  may 
be  detected  by  the  brown-ring  test.  This  test  is  made  by  adding  a 
small  amount  of  freshly  prepared  ferrous  sulfate  solution  to  the  solu- 
tion suspected  of  containing  a  nitrate.  Concentrated  sulfuric  acid  is 
then  carefully  poured  down  the  side  of  the  test  tube  in  such  a  way 
that  it  collects  at  the  bottom  without  mixing  with  the  solution.  If 
a  brown  layer  forms  between  the  heavy  sulfuric  acid  and  the  solu- 
tion floating  on  top,  a  nitrate  is  present.  The  sulfuric  acid  pro- 
duces nitric  acid  by  reacting  with  the  nitrate,  and  the  nitric  acid 
in  turn  reacts  with  the  ferrous  sulfate,  forming  a  brown  compound, 
FeSO4  •  NO. 

Gunpowder,  the  earliest  explosive.  The  inactive  element  nitrogen 
does  not  unite  easily  with  other  elements.  And  after  it  does,  the 
unions  so  produced  are  very  unstable.  In  fact,  such  unions  are  so 
extremely  unstable  that  on  the  slightest  provocation  *  the  nitrogen 
breaks  away  with  a  bang!  Most  explosives,  except  those  based  on 
nuclear  fission,  depend  upon  this  fact. 

Nearly  all  nonfission  explosives  contain  either  nitrate  (NO3) ,  or 
nitro  (NO2) ,  radicals.  In  addition,  some  explosives  contain  ammo- 
nium radicals,  NH4.  When  compounds  that  contain  nitrate  or  nitro 
radicals  are  mixed  with  other  compounds  that  can  easily  use  the  oxy- 
gen of  these  unstable  radicals,  an  explosive  is  the  result.  In  certain 
cases,  the  compound  containing  the  nitrate  or  nitro  radicals  actu- 
ally supplies  the  means  of  its  own  destruction  by  furnishing  the  ele- 
ments that  can  use  the  oxygen  readily.  When  something,  such  as  a 
shock,  starts  the  reaction,  the  unstable  nitrate  or  nitro  radicals  re- 
lease their  oxygen  for  combination  with  other  elements  and  liberate 
free  nitrogen  gas.  Nearly  always,  most  of  the  other  products  of  the 
reaction  are  gases  also,  and  because  of  the  high  temperatures  pro- 
duced, terrific  pressures  result. 

The  earliest  explosive  was  made  thousands  of  years  ago  by 
the  Chinese.  How  the  black  powder  came  to  be  produced  is  not 
known,  but  it  was  made  then  about  as  it  is  now,  by  mixing  approxi- 
mately 15  parts  by  weight  of  potassium  nitrate,  KNO3,  with  three 
parts  of  sulfur,  and  two  parts  of  powdered  charcoal.  The  resulting 


black  powder,  or  gunpowder,  explodes  with  terrific  violence,  much 
lire,  and  great  quantities  of  acrid  smoke.  The  reaction  produced  is: 

2KNO3  +  3C  +  S  -»  K2S  +  Ns  t  +  3CO2  T 

The  Chinese  used  gunpowder  in  producing  firecrackers  and  other 
kinds  of  fireworks  for  use  in  ceremonies  and  celebrations.  But  some- 
time within  the  past  few  centuries,  no  one  knows  for  sure  just  when, 
gunpowder  was  turned  to  military  uses,  and,  like  so  many  other  sci- 
entific discoveries,  soon  created  a  revolution.  The  foot  soldier  with 
his  primitive  musket  quickly  replaced  the  heavily-armored  and  be- 
pl  umed  knight  with  his  lance  and  shield.  Turrets  and  thick  stone 
walls  were  no  defense  against  powerful  cannons,  and  the  picturesque 
castle  of  the  Middle  Ages  became  obsolete. 

Nitrocellulose,  nitroglycerin,  and  some  detonators.  Today,  gun- 
powder is  considered  a  relatively  "tame"  explosive.  Since  its  inven- 
tion, chemists  have  produced  many  kinds  of  explosives  chiefly  by 
nitrating,  or  adding  NCX  or  NO,  radicals  to,  such  substances  as  cot- 
ton, glycerin,  sugar,  starch,  and  other  organic  compounds.  Nitrocel- 
lulose, or  guncotton,  was  produced  by  Schonbein  (shun'bln)  in  1846 
by  nitrating  cotton  with  a  mixture  of  nitric  and  su  If  uric  acids.  Nitro- 
glycerin, made  by  nitrating  glycerin,  a  common  byproduct  of  the 
manufacture  of  soap,  was  produced  in  1847.  Both  of  these  compounds 
are  more  powerful  explosives  than  gunpowder,  but  both  are  much 
more  sensitive  to  shock  and,  hence,  explode  much  more  easily. 

In  1888,  Alfred  Nobel,  a  Swedish  inventor,  produced  dynamite 
by  absorbing  nitroglycerin  in  a  fine  clay,  or  diatomaceous  earth  (see 
page  500)  .  Dynamite  is  much  less  sensitive  to  shock  and,  hence, 

Standard  Oil  Company  (\'.J.) 

A  workman  inserts  a  dynam 
cartridge   into   a    drilled    h 
during  the  construction  of 
underground  pipeline. 



can  be  used  with  greater  safety  than  nitroglycerin.  American  dy- 
namite usually  consists  of  nitroglycerin  absorbed  in  a  wood  meal 
that  resembles  fine  sawdust.  The  substance  is  packed  in  sticks  com- 
posed of  parchment  paper.  So  overcome  was  Nobel  by  the  possible 
uses  of  his  invention  that  he  dedicated  part  of  its  profits  to  the  es- 
tablishment of  the  Nobel  Peace  prize  for  outstanding  contributions 
to  the  peace  of  the  nations  of  the  world,  and  of  the  Nobel  prizes  for 
outstanding  contributions  to  research  in  physics,  chemistry,  medi- 
cine and  physiology,  and  literature.  Peacetime  uses  of  explosives  in- 
clude mining,  building  dams,  and  other  construction  work. 

Explosives  are  set  off,  or  detonated,  by  means  of  a  shock  produced 
by  even  more  unstable,  and  thus  more  sensitive,  compounds  of  ni- 
trogen called  detonators.  Fulminate  of  mercury,  Hg  (CNO)  2,  a 
widely  used  detonator,  is  1 1  times  as  sensitive  to  shock  as  trinitro- 
toluene, or  TNT,  and  twice  as  sensitive  as  nitrocellulose.  Lead  azide, 
Pb  (N3)  2,  another  detonator,  is  half  as  sensitive  to  shock  as  fulminate 
of  mercury.  Such  substances  are  used  in  making  the  caps  and  other 
devices  with  which  explosives  are  set  off. 

Nitric  acid  has  many  peacetime  uses.  Aside  from  the  production 
of  explosives,  a  major  industry  even  in  peacetime,  the  chief  use  of 
nitric  acid  is  in  the  production  of  nitrates  of  organic  compounds, 
such  as  nitrocellulose  and  nitrobenzene.  Nitrocellulose  is  used  in 
making  some  photographic  films,  and  many  quick-drying  lacquers 
and  enamels,  especially  for  the  automobile  industry.  Nitrocellulose 
is  used  also  in  the  production  of  many  kinds  of  artificial  leathers. 
Nitrobenzene  is  the  basic  raw  material  of  the  aniline,  or  coal-tar 
dye,  industry. 

Nitric  acid  also  furnishes  the  oxides  of  nitrogen  required  in  the 
chamber  process  for  the  manufacture  of  sulfuric  acid  (see  page  311). 
Nitrates  for  fertilizers  and  metallic  nitrates  are  made  from  nitric 
acid.  Sodium  nitrate  and  ammonium  nitrate  are  the  chief  fertilizers 
produced.  Strontium  nitrate  is  one  of  the  chemicals  used  in  pyro- 
technics, the  production  of  fireworks,  which  consist  mainly  of  flares 


Use  by  approximate  percentage 


and  shells  that  give  off  flames  and  smokes  of  various  colors.  A  one 
percent  silver  nitrate  solution  is  put  in  the  eyes  of  newborn  babies 
to  prevent  infection  that  may  lead  to  blindness. 

Nitric  acid  plays  an  important  role  in  the  pickling  of  steel,  in  the 
etching  of  engravers'  plates  for  printing,  and  in  the  manufacture  of 
the  arsenate  insecticides  (chiefly  lead  and  calcium  arsenates)  so 
widely  used  against  the  boll  weevil,  and  in  the  spraying  of  fruit 
trees  (see  page  453) . 

The  nitrogen  situation  today  represents  a  chemical  revolution. 
It  should  be  apparent  already  that  the  production  of  nitrogen  com- 
pounds is  a  basic  industry.  The  world's  normal  consumption  of  man- 
ufactured nitrogen  compounds  is  many  millions  of  tons  annually. 
At  one  time,  sodium  nitrate  from  Chile  and  nitrogen  compounds  ob- 
tained as  byproducts  from  the  coal  industry  and  slaughterhouses 
were  the  only  sources  of  nitrogen  compounds.  Today,  the  total  an- 
nual amount  of  fixed  nitrogen  produced  by  chemical  methods  is 
many  times  as  great  as  the  total  annual  consumption  of  both  Chilean 
nitrate  and  all  byproduct  nitrogen  taken  together.  The  story  of  the 
nitrogen  industry  bears  testimony  to  the  widespread  development 
of  synthetic  chemistry.  It  has  changed  the  economic  life  of  millions. 

There  are  six  oxides  of  nitrogen.  In  addition  to  the  two  oxides 
of  nitrogen  already  discussed,  four  other  oxides  are  known.  The  com- 
plete list  is:  nitrous  oxide,  N2O;  nitric  oxide,  NO;  nitrogen  diox- 
ide, NO2;  nitrogen  trioxide,  N2O^  nitrogen  tetroxide,  N.X)4,  a  pow- 
erful oxidizing  agent;  and  nitrogen  pentoxide,  N2Or,.  They  illustrate 
the  law  ot  multiple  proportions  and  the  fact  that  nitrogen  has  sev- 
eral different  valences.  Why? 

When  nitrogen  trioxide,  N2Oy,  is  added  to  water,  nitrous  acid,  a 
very  unstable  acid,  is  formed. 

N2O3  +  H2O  ->  2HNO2 

Similarly,  the  addition  of  water  to  nitrogen  pentoxide,  N2O5,  pro- 
duces nitric  acid. 

N205  +  H20  ->  2HN03 

These  two  gases  may  therefore  be  said  to  be  the  anhydrides  (with- 
out water)  of  nitrous  and  nitric  acids,  respectively.  An  acid  anhy- 
dride is  an  oxide  whose  water  solution  is  an  acid. 

Nitrous  oxide,  or  laughing  gas.  Priestley  was  the  first  to  produce 
nitrous  oxide,  N2O,  a  colorless,  heavy  gas,  slightly  sweetish  in  odor 
and  somewhat  soluble  in  water.  This  was  about  two  years  before  he 

A  modern  hospital  operating 
room.  Suspended  from  the 
table  in  the  left  foreground 
are  tanks  of  pure  oxygen, 
carbon  dioxide,  nitrous  oxide 
and  other  gases  ready  for  im- 
mediate administration  when 

Ohio  Chemical  and  Surgical  hqu, 

discovered  oxygen.  He  learned  that  it  supported  the  burning  of  a 
candle  better  than  did  ordinary  air.  It  decomposes  rather  easily  into 
oxygen  and  nitrogen.  Just  before  the  close  of  the  eighteenth  century, 
Humphry  Davy  achieved  fame  overnight  by  his  discovery  of  the 
physiological  effects  of  this  gas.  He  breathed  four  gallons  of  it  and 
noticed  its  power  to  produce  a  peculiar  intoxication,  which  included 
laughing.  The  poet  Samuel  Coleridge,  as  well  as  other  distinguished 
persons,  came  to  Davy's  London  laboratory  to  experience  the  thrill 
of  inhaling  this  gas.  Nitrous  oxide  is  still  prepared  as  Davy  made  it, 
by  heating  ammonium  nitrate. 

NH4N03  ->  2H20  +  N2O 

In  1842  ether  was  used  as  the  first  anesthetic  in  surgery  by  Dr. 
Crawford  W.  Long,  a  country  doctor  of  Georgia.  William  Morton's 
use  of  ether  at  the  Massachusetts  General  Hospital  in  1846  intro- 
duced this  anesthetic  to  the  medical  world.  Two  years  earlier,  Dr. 
Horace  Wells,  a  dentist  of  Hartford,  Connecticut,  had  one  of  his 
teeth  extracted  after  he  had  anesthetized  himself  with  nitrous  oxide. 

Today  nitrous  oxide  is  still  used  as  an  anesthetic  in  many  opera- 
tions, especially  those  of  dentistry.  It  is  usually  mixed  with  about 
25  percent  oxygen  and,  in  cases  of  more  serious  operations,  with 
ether.  This  mixture  of  nitrous  oxide  and  oxygen  can  be  breathed 
for  a  considerable  period  without  harmful  effects  on  the  circulatory 



system  or  on  vital  organs.  Small  amounts  of  nitrous  oxide  are  used 
in  preserving  perishable  foods  and  liquids.  Easily  liquefied,  it  is  sold 
in  cylinders.  It  is  used  to  eject  whipped  cream  at  soda  fountains. 


Conant,  James  Bryant.  The  Overthrow  of  the  Phlogiston 
Theory.  Case  2  of  the  Harvard  Case  Histories  in  Experimental 
Science.  Harvard  University  Press,  Cambridge,  Mass.,  1950. 
Gives  an  excellent  account  of  Priestley's  confusion  between 
oxygen  and  nitrous  oxide. 

Haynes,  William.  This  Chemical  Age,  pp.  78-94.  Alfred  A. 
Knopf,  New  York,  1942.  A  discussion  on  explosives  and  their 
relation  to  the  dye  industry,  entitled  "Mars:  Chemical  Dic- 

Slosson,  Edwin  E.  Creative  Chemistry,  pp.  37-59.  D.  Apple- 
ton-Century  Co.,  New  York,  1920.  A  very  readable  account  of 
nitrogen  and  nitrogen  compounds  in  relation  to  plants. 


Group  A 

1.  Explain  the  arc  process  of  making  HNO3. 

2.  Write  equations  showing  two  methods  of  preparing  NO. 

3.  State  one  chemical  and  three  physical  properties  of  NO. 

4.  Make  a  diagram  showing  the  laboratory  preparation 
of  NO. 

5.  (a)  What  happens  when  NO  comes  in  contact  with 
air?  (b)  Explain. 

6.  Under  what  conditions  does  NO0  change  into  N.,O4? 


7.  Write  the  equation  that  is  the  basis  of  the  Ostwald  proc- 
ess for  the  synthesis  of  HNO3. 

8.  (a)  What  natural  phenomenon  results  in  the  formation 
of  certain  oxides  of  nitrogen?  (b)  Explain. 

9.  (a)  What  is  nitrogen  fixation?  (b)  Illustrate. 

10.  Make  a  diagram  showing  the  laboratory  preparation  of 

11.  When  a  mixture  of  NaNO3  and  H2SO4  is  heated  in  a 
retort,  HNO3  is  formed  a  little  at  a  time,    (a)  What  four 


substances  are  in  the  retort?   (b)  Which  is  removed  by  heat? 
Why?  (c)  Why  do  the  other  substances  remain? 

12.  (a)  What  property  of  HNO3  makes  it  possible  to  pre- 
pare the  acid  by  the  laboratory  method?  (b)  Why  could  not 
HC1  be  used  instead  of  H2SO4? 

13.  Write  the  equation  for  the  laboratory  preparation  of 

14.  State  four  chemical  properties  of  HNO3. 

15.  (a)  How  is  aqua  regia  prepared?    (b)  Its  power  to  dis- 
solve gold  results  from  what  property?   (c)  What  property  of 
HNO3  is  shown? 

16.  (a)  Why  does  HNO3  appear  to  be  yellow  when  prepared 
in  the  laboratory?  (b)  What  property  of  HNOrf  does  this  color 

17.  (a)  What   oxide   of  nitrogen   is   always   formed   when 
HNO3   decomposes?    (b)  What  oxide  of  nitrogen   is   always 
formed  when  HNO3  acts  as  an  oxidizing  agent? 

18.  How  would  you  test  for  the  presence  of  the  nitrate  ion? 

19.  What  are  the  three  principal  uses  of  HNO/ 

20.  (a)  Most  explosives  are  based  upon  what  fact?  (b)  What 
is  a  detonator? 

21.  (a)  What  is  gunpowder?   (b)  dynamite? 

22.  Copy  and  complete  the  following:   The  six  oxides  of 
nitrogen  illustrate  the  law  of  ....  The  anhydride  of  HNO3 
is  .  .  .  ,  and  the  anhydride  of  HNO2  is  ....  Another  name  for 
N2O  is  ....  N2O  was  first  used  as  ...  by  Dr.  Horace  Wells. 
N2O  is  prepared  by  heating  .... 

23.  120  g.  of  NO  are  obtained  by  the  action  of  Cu  on 
HNO3.  How  many  grams  of  Cu  (NO3)  2  are  formed? 

24.  How   do   you   explain   the    fact   that    the   reaction   of 
N2  -(-  O2  —  »  2NO,  although  known  for  more  than  a  century, 
could  not  be  used  in  the  preparation  of  HNO.  until  rather 

I  ... 

25.  Why  is  N2O  a  better  supporter  of  combustion  than  NO? 

26.  (a)  How  does  nature  restore  some  of  the  nitrogen  com- 
pounds taken  from  the  soil  by  growing  crops?    (b)  What  is 
"green  manure"? 


27.  In  1898,  Sir  William  Crookes,  one  of  England's  most 
eminent  chemists,  predicted  a  future  world  famine  caused  by 
exhaustion  of  nitrogen  compound  fertilizers.  Why  has  his  pre- 
diction not  materialized? 

28.  Make  a  table  showing  the  properties  of  the  six  oxides  of 

29.  How  would  you  separate  Au  from  Cu  in  a  copper-gold 

Group  B 

30.  Why  cannot  HNO3  be  used  in  preparing  H2?  Answer 
this  question  in  the  light  of  the  electron  theory. 

31.  (a)  Using  the  electron  theory,  explain  how  HNO3  oxi- 
dizes Cu.  By  inspection  of  the  equation,  state  how  many  atoms 
of  Cu  are  oxidized,    (b)  How  many  atoms  of  nitrogen  are  re- 
duced?  (c)  How  does  the  total  number  of  electrons  lost  com- 
pare with  the  total  number  gained? 

32.  How  could  you  determine  experimentally  whether  a 
gas  contained  a  high  percentage  of  N2O4  and  a  small  amount 
of  NO2,  or  was  composed  almost  entirely  of  NO,? 

33.  There  is  less  oxygen  in  nitrous  oxide   than  in  nitric 
oxide,  (a)  Which  would  support  burning  better?  (b)  Explain. 

34.  Nitrates  are  unstable  in  the  presence  of  heat.  What 
product  would  you  expect  Cu  (NO3)  2  to  yield  when  heated? 


1.  At  Muscle  Shoals,  Alabama,  there  is  a  huge  plant  for 
generating  electric  power.  This  plant  is  capable  of  producing 
thousands  of  tons  of  nitrogen  compounds  a  year  by  fixation. 
The  U.S.  government  spent  many  millions  of  dollars  on  this 
project.  What  is  the  present  state  of  this  plant?  Consult  your 
teacher  of  history  or  economics  or  write  to  your  Congressman, 
Senator,  or  the  Tennessee  Valley  Authority  at  Knoxville,  Ten- 
nessee. What  do  you  think  should  be  done?  Prepare  a  debate 
or  discussion  of  this  question. 

2.  Write  a  report  comparing  the  changes  brought  about  by 
the  introduction  of  gunpowder  into  warfare  with  the  changes 
that  the  A-  and  H-bombs  may  produce. 

3.  If  available  at  this  time,  obtain  a  fresh  sample  of  some 
legume  such  as  peas,  beans,  alfalfa,  or  soybeans.  Examine  its 
roots  for  nodules  of  nitrogen-fixing  bacteria.  Exhibit  to  class 
and  explain  its  function. 




.  .  .  In  1858  the  atomic  theory  of 
Da  I  ton  was  just  50  years  old.  Stu- 
dents at  this  time  were  generally  un- 
familiar with  the  word  molecule,  for 
chemists  spoke  as  complacently 
about  an  atom  of  water  as  about  an 
atom  of  oxygen.  For  the  most  part, 
also,  they  had  never  heard  of  Amedeo 
Avogadro.  William  Tilden,  1921 

The  battle  over  the  molecules  of  Avogadro.  In  1860,  chemical 
science  was  in  a  turmoil  caused  by  a  misunderstanding  of  the  terms 
atom  and  molecule.  Chemists  had  spoken  of  "atoms  of  water,"  which 
is  a  compound,  in  the  same  way  in  which  they  mentioned  atoms  of 
hydrogen,  which  is  an  element.  Some  used  the  term  compound 
atoms.  So  great  was  the  confusion  that  finally  a  congress  of  chem- 
ists was  called  at  Karlsruhe  to  decide  when  to  use  the  term  atom  and 
when  to  use  the  term  molecule,  which  up  to  that  time  had  been  used 

Among  the  brilliant  men  who  gathered  at  Karlsruhe  from  all  parts 
of  the  world  was  a  bearded  young  Italian,  Cannizzaro  (kan-net- 
sa'ro)  .  He  had  come  to  champion  the  use  of  the  term  molecule  in 
the  sense  that  it  was  used  in  1811  by  Avogadro  (a-vo-ga'dro) ,  a  mod- 
est professor  of  chemistry.  According  to  Avogadro,  a  molecule  is  the 
smallest  part  of  either  an  element  or  a  compound  zvhich  has  the  prop- 
erties of  that  substance. 

Avogadro  had  reached  this  new  meaning  of  a  molecule  from  his 
study  of  the  behavior  of  gases.  He  had  died  four  years  before  the  con- 
gress was  held,  but  Cannizzaro  championed  his  ideas  so  successfully 
that  they  were  finally  accepted  by  the  congress. 




What  Boyle  and  Charles  discovered  about  the  behavior  of  gases. 
In  order  to  understand  the  significance  of  Avogadro's  contribution 
to  chemistry,  it  is  necessary  to  trace  the  story  of  the  study  of  gases 
after  1660.  In  that  year  Robert  Boyle,  revered  by  Englishmen  as 
the  father  of  modern  chemistry,  discovered  that  if  the  pressure  on 
a  gas  is  doubled  —  for  example,  increased  from  15  pounds  to  30 
pounds  per  square  inch  —  its  volume  is  decreased  one-half.  Further- 
more, he  found  that  this  relationship  between  pressure  and  volume 
does  not  depend  upon  the  nature  of  the  gas;  it  is  true  for  all  gases. 
Thus  we  have  Boyle's  law:  The  volume  of  a  gas  varies  inversely  as 
the  pressure  exerted  upon  it  if  the  temperature  remains  constant. 

In  1785,  Charles,  a  French  scientist,  noted  that  under  constant 
pressure  the  volume  of  a  gas  increases  ^^  of  its  volume  at  0°C.  for 
each  centigrade  degree  of  rise  in  temperature;  in  other  words,  if  the 
absolute  temperature  (see  page  644)  of  a  gas  increases  from  273°  to 
546°  (2  X  273°)  its  volume  doubles  also.  This  he  found  true  of  any 
gas.  Thus  we  have  Charles'  law:  The  volume  of  a  gas  varies  directly 
as  the  absolute  (A)  temperature,  if  the  pressure  on  the  gas  remains 

Although  gases  act  according  to  Boyle's  and  Charles'  laws  under 
ordinary  temperatures,  they  do  not  do  so  at  very  high  pressures  or 
very  low  temperatures.  However,  a  discussion  of  the  way  in  which 
gases  act  at  low  temperatures  and  under  high  pressures  is  beyond  the 
scope  of  an  introductory  course  in  chemistry. 

Gay-Lussac  finds  the  law  of  combining  volumes  of  gases.  The 
next  important  advance  came  23  years  later.  Gay-Lussac,  the  French- 
man who  collected  air  samples  in  a  balloon  high  over  Paris,  had 
long  been  interested  in  the  study  of  gas  volumes.  He  knew  that  one 
volume  of  nitrogen  unites  with  one  volume  of  oxygen,  forming  two 
volumes  of  nitric  oxide.  Besides,  when  he  repeated  the  experiments 
of  Cavendish  and  Lavoisier,  he  found  that  two  volumes  ot  hydrogen 
unite  with  one  volume  of  oxygen,  forming  two  volumes  of  water  va- 
por. This  was  not  a  new  discovery,  but  Gay-Lussac  suspected  that 
"other  gases  might  also  combine  in  simple  ratios." 

plunger  I 



Fig.  63.  A  demonstration  of 
the  principles  of  Boyle's 


Resuming  his  researches,  Gay-Lussac  discovered  that  one  .volume 
of  hydrogen  chloride  gas  when  brought  in  contact  with  one  volume 
of  ammonia  gas  yielded  a  white  powder,  with  no  residue  of  either 
gas.  The  two  gases  had  combined  volume  for  volume.  Furthermore, 
he  had  read  that  one  volume  of  nitrogen  combined  with  exactly 
three  volumes  of  hydrogen,  forming  exactly  two  volumes  of  ammo- 
nia gas.  This  was  an  arithmetical  simplicity  of  remarkable  signifi- 
cance. Fractions  of  volumes  of  gases  were  not  involved. 

1  vol.  nitrogen  +  1  vol.  oxygen  —-»  2  vol.  nitric  oxide 
1  vol.  oxygen  +  2  vol.  hydrogen  — »  2  vol.  water  vapor 
1  vol.  HC1  +  1  vol.  NH3  -»  NH4C1  (a  solid) 

1  vol.  nitrogen  -f  3  vol.  hydrogen  — *  2  vol.  ammonia  gas 

On  the  last  day  of  the  year  1808,  Gay-Lussac  formulated  from  these 
observations  a  law,  which  bears  his  name.  Gay-Lussac's  law  states 
that  the  relation  between  the  combining  volumes  of  gases  and  the 
volumes  of  their  products  (if  they,  too,  are  gases)  may  be  expressed 
in  small  whole  numbers.  Why  this  regularity?  On  the  basis  of  Dai- 
ton's  atomic  theory,  chemists  could  not  explain  this  law. 

When  Dal  ton  was  faced  with  this  fact,  he  refused  to  accept  Gay- 
Lussac's  law.  "The  truth  is,"  Dalton  maintained,  "that  gases  do  not 
unite  in  equal  or  exact  proportions  in  any  one  instance.  When  they 
appear  to  do  so,  it  is  owing  to  the  inaccuracy  of  our  experiments." 
Later,  however,  after  further  experimentation  and  study,  Dalton 
accepted  Gay-Lussac's  law.  His  acceptance  of  Gay-Lussac's  law,  when 
experimental  evidence  pointed  toward  its  accuracy,  reveals  Dalton 
as  a  true  scientist. 

Is  the  sum  of  2  and  1  always  3?  The  law  of  Gay-Lussac  in  par- 
ticular, and  the  laws  of  Boyle  and  Charles  to  a  lesser  degree,  sug- 
gested a  number  of  interesting  problems  to  Avogadro's  inquiring 
mind.  Why,  for  example,  is  the  behavior  of  gases  so  uniform  under 
changing  temperature,  while  the  behavior  of  solids  and  liquids  is 
so  variable?  Why  do  gases  combine  in  simple  ratios  by  volume?  Fur- 
thermore, why,  with  respect  to  gases,  is  not  the  sum  of  two  and  one 
always  three?  For  example,  why  do  two  volumes  of  hydrogen  unite 
with  one  volume  of  oxygen,  making  two  volumes  of  water  vapor;  and, 
similarly,  three  volumes  of  hydrogen  combine  with  one  volume  of 
nitrogen,  making  two  volumes  of  ammonia  gas? 

Avogadro  continues  his  work  with  gases.  Avogadro  tried  to  answer 
these  questions.  In  reference  to  the  last  question,  perhaps  he  thought: 
it  might  be  that  equal  volumes  of  gases  contain  the  same  number  of 


But  according  to  this  idea,  one  volume  of  oxygen  ought  to  combine 
with  one  volume  of  nitrogen,  forming  one  volume  of  nitric  oxide: 

N  +  O  -»  NO 

whereas,  according  to  actual  experiment,  two  volumes  of  nitric  oxide 
are  formed.  Something  was  wrong.  But  wait!  Suppose  the  molecule 
of  nitrogen  gas  contains  two  atoms  instead  of  one,  that  is,  is  N2  and 
not  N,  and  similarly  the  oxygen  molecule  is  O2  and  not  O,  what  then? 
According  to  this  idea,  the  equation  would  be: 

N2  +  O2  ->  2NO 

and  the  conditions  necessary  for  the  formation  of  two  volumes  of  NO 
would  be  fulfilled. 

How  would  this  idea  work  out  in  other  cases,  for  example,  in  the 
formation  of  water?  If  the  molecule  of  hydrogen  contains  two  atoms, 
like  the  molecule  of  oxygen,  we  should  have  the  equation, 

2H2  +  O2  ->  2H2O 

or  stated  in  other  words,  two  volumes  of  hydrogen  unite  with  1  vol- 
ume of  oxygen,  forming  two  volumes  of  water  vapor.  This  agrees 
with  actual  measurements  of  the  volumes. 

Apparently  Avogadro  was  on  the  right  track.  It  remained  for  him 
only  to  test  his  hypothesis  further  by  means  of  other  gas  combina- 
tions to  be  able  to  show  that,  assuming  the  molecules  of  elementary 
gases  to  be  composed  of  two  atoms  each,  the  volumes  corresponded 
with  the  equation  as  he  had  calculated. 

This  he  actually  did  and  finally  was  able  to  establish  the  accuracy 
of  his  hypothesis,  that  all  gases  behave  alike,  because  equal  volumes 
of  all  gases  under  the  same  conditions  of  temperature  and  pressure 
are  composed  of  the  same  number  of  molecules  (Avogadro's  hypoth- 
esis) .  To  this  professor  from  Turin,  elementary  gases,  such  as  hydro- 
gen, oxygen,  nitrogen,  and  chlorine,  consist  normally  of  molecules 
each  composed  of  two  atoms  instead  of  one,  as  Dalton  and  the  rest 
of  the  world  had  supposed.  Incidentally,  Avogadro's  hypothesis  rec- 
onciled the  atomic  theory  of  Dalton  with  Gay-Lussac's  law.  How? 

Of  what  value  to  chemistry  was  Avogadro's  hypothesis?  What 
evidence  did  Avogadro  have  to  back  up  so  bold  an  hypothesis?  He 
could  not  verify  it  experimentally.  No  balance  was  sensitive  enough 
to  weigh  a  molecule.  It  would  take  billions  of  these  tiny  particles 
to  turn  the  scales  of  even  the  most  sensitive  balance.  He  surely  had 
not  looked  into  the  molecules  of  matter  and  detected  the  twin 


arrangement  of  atoms,  for  it  would  take  many  millions  of  mole- 
cules placed  side  by  side  to  make  a  line  one  inch  long.  Only  in  re- 
cent years  have  methods  been  developed  which  can  make  such  a 
tiny  particle  visible. 

The  only  evidence  Avogadro  had  was  that  of  clear,  accurate  rea- 
soning and  his  own  creative  imagination.  However,  this  evidence 
was  strong  enough  to  clear  the  air  and  allow  chemistry  to  advance. 
The  particles  of  elementary  gases  were  henceforth  considered  to  be 
diatomic,  that  is,  composed  of  two  atoms  to  the  molecule.  (Later, 
by  other  methods,  the  inert  gases  of  the  atmosphere  were  shown  to 
contain  only  one  atom  to  the  molecule.)  Atomic  weights  and  molecu- 
lar weights  were  thus  clearly  differentiated.  New  methods  were  made 
possible  for  determining  the  molecular  weights  of  gases  and,  from 
these,  their  atomic  weights  also. 

How  Avogadro's  hypothesis  was  actually  verified.  Since  the  time 
of  Avogadro,  new  apparatus  and  new  methods  have  been  devised  for 
verifying  his  prophetic  statement.  A  number  of  scientists,  Millikan 
and  Perrin  (pe-raN')  among  them,  determined  by  experiment  the 
number  of  molecules  in  a  given  volume  of  gas.  They  found  that  the 
number  of  molecules  in  two  grams  of  hydrogen  gas  (a  gram-molecu- 
lar weight) ,  for  example,  is  approximately  602,000,000,000,000,000,- 
000,000  (602  sextillion) .  This  number,  usually  written  6.02  X  1023, 
is  now  called  Avogadro's  number.  Approximately  this  number  of 
molecules  is  known  to  be  present  in  equal  volumes  (22.4  liters)  of 
all  gases  and  vapors  under  the  same  conditions  of  temperature  and 
pressure.  This  is  no  idle  guess. 

Perrin  and  Millikan,  both  Nobel  prize  winners  in  physics,  main- 
tained that  we  can  count  the  number  of  molecules  in  a  small  volume 
of  a  gas  with  as  much  accuracy  as  we  can  determine  the  population 
of  a  city  such  as  New  York.  Avogadro's  hypothesis  has  now  taken  its 
place  as  one  of  the  laws  of  chemistry. 

Then  came  another  remarkable  verification.  Irving  Langmuir 
(lang'mur) ,  another  Nobel  laureate  in  chemistry,  succeeded  in  break- 
ing up  the  molecules  of  hydrogen  gas.  As  a  result  of  his  experiments 
he  found  that  hydrogen  gas  is  made  up  of  molecules  each  of  which 
consists  of  two  atoms.  Langmuir  made  use  of  this  discovery  when  he 
invented  the  atomic-hydrogen  torch. 

Principle  of  the  atomic-hydrogen  torch.  In  an  atomic-hydrogen 
torch,  hydrogen  gas  is  passed  through  an  electric  arc  produced  be- 
tween electrodes  made  of  wolfram.  The  heat  of  the  electric  arc  splits 
the  hydrogen  molecule  into  hydrogen  atoms.  Immediately  after  pass- 
ing through  the  arc,  the  atoms  reunite,  forming  hydrogen  molecules, 

General  Electric  Company 

Repairing  worn  parts  of  vacuum  cleaners  with  an  atomic-hydrogen  arc  welder. 

which  are  oxidized,  forming  water.  Atomic  hydrogen  cannot  be 
stored.  ^ 

H2  *=>  H  +  H 

All  the  energy  absorbed  from  the  electric  arc  in  splitting  the  mole- 
cule is  liberated  when  the  atoms  of  hydrogen  reunite.  This  energy, 
as  heat,  added  to  the  heat  normally  generated  when  hydrogen  burns, 
produces  a  temperature  between  4000°C.  and  5000°C.  (The  oxy- 
acetylene  torch  gives  a  temperature  of  about  3300 °C.) 

The  atomic-hydrogen  torch  is  used  for  cutting  and  welding  metals. 
It  has  the  advantage  of  protecting  the  object  against  oxidation,  since 
the  jet  of  burning  hydrogen  is  always  surrounded  by  hydrogen,  a 
reducing  agent. 

Proof  that  the  molecule  of  hydrogen  contains  two  atoms.  The- 
oretically, we  can  prove  the  formula  of  hydrogen  gas  to  be  H,  as 
follows:  (1)  From  experiments,  we  know  that  one  volume  of  hydro- 
gen unites  with  one  volume  of  chlorine,  yielding  two  volumes  of  hy- 
drogen chloride  gas.  (2)  According  to  Avogadro's  law,  equal  vol- 
umes of  all  gases  contain  the  same  number  of  molecules.  Conversely, 
equal  numbers  of  molecules  of  gases  occupy  equal  volumes.  Hence, 



one  molecule  of  hydrogen  and  one  molecule  of  chlorine  occupy  ,equal 
volumes,  and  two  molecules  of  HC1  gas  occupy  twice  this  volume. 
We  may  represent  this  graphically  as  follows: 

1  volume      4-     1  volume         +~       2  volumes (by  experiment) 

1  molecule    4-   1  molecule       ^     \2  molecules/- --(by  Avogadro's 

of  hydrogen  \  of  HCI  /  law) 

must  contain 1  must  contain 

' +~  at  leasF*2"atoms  of  hydrogen 

(3)  One  of  the  HCI  molecules  must  contain  at  least  one  atom  of  hy- 
drogen, since  fractions  of  atoms  do  not  exist.  (4)  Since  we  have  two 
HCI  molecules,  we  must  have  at  least  two  atoms  of  hydrogen  which 
can  have  come  from  only  the  one  molecule  of  hydrogen. 


1.  We  know  that  1  vol.  of  nitrogen  unites  with  1  vol.  of 
oxygen,  forming  2  vol.  of  nitric  oxide.  Prove  that  the  mole- 
cule of  nitrogen  contains  two  atoms. 

2.  One  vol.  of  hydrogen  unites  with  1  vol.  of  bromine  vapor, 
forming  2  vol.  of  hydrogen  bromide  gas.  Show  that  the  mole- 
cule of  bromine  vapor  contains  two  atoms. 

3.  One  vol.  of  oxygen  unites  with  2  vol.  of  hydrogen,  form- 
ing 2  vol.  of  water  vapor.  Show  that  the  formula  for  oxygen 
gas  is  O2. 

4.  Prove  the  structure  of  the  hydrogen  molecule  from  the 
fact  that  1  vol.  of  nitrogen  unites  with  3  vol.  of  hydrogen, 
forming  2  vol.  of  ammonia  gas. 

The  gram-molecular  volume  of  a  gas  or  vapor.  You  have  learned 
that  a  chemical  formula  may  represent  one  molecule  of  a  compound; 
one  molecular  weight  of  a  compound;  and  also  one  gram-molecular 
weight,  or  mole,  ot  a  compound.  For  example,  CO2  may  stand  for 
one  molecule  of  carbon  dioxide,  for  the  molecular  weight  of  car- 
bon dioxide  (44) ,  or  for  one  mole  of  carbon  dioxide  (44  grams) .  If 
a  coefficient  appears  in  front  of  a  formula,  it  represents  a  definite 
number  of  molecules,  molecular  weights,  or  moles.  Thus  2CO2  stands 
for  two  molecules  of  the  gas,  two  molecular  weights  (88) ,  or  two 
moles  (88  grams) .  A  formula  has  an  additional  meaning  which  is 
important  in  many  chemical  calculations. 


In  dealing  with  a  gas  or  vapor,  it  is  often  necessary  to  know  the 
volume  that  a  quantity  of  it  occupies.  The  unit  of  measurement  of 
gas  volumes  is  the  volume  occupied  by  one  mole  (abbreviated  M)  . 
This  is  called  the  gram-molecular  volume  (V)  .  Study  has  shown  that 
the  gram-molecular  volume  is  the  same  for  all  gases.  Under  standard 
conditions  of  temperature  and  pressure,  one  M  of  any  gas  or  vapor  oc- 
cupies 22.4  liters.  This  may  be  demonstrated  by  the  experimental 
process  of  weighing  a  given  volume  of  any  gas.  For  example,  one 
liter  of  hydrogen  weighs  0.08987  gram.  Therefore  22.4  liters  weigh 
2.016  grams,  which  is  the  gram-molecular  weight  (mole)  of  hydrogen. 

Since  one  mole  of  any  gas  occupies  22.4  liters,  we  may  use  the 
formula  of  the  gas  to  represent  its  gram-molecular  volume.  For  exam- 
ple, CO2,  which  represents  one  mole  of  carbon  dioxide,  also  repre- 
sents one  gram-molecular  volume  (V)  of  carbon  dioxide;  NH3  rep- 
resents one  M  of  ammonia  and  also  one  V  of  ammonia.  In  each  case, 
V  =  22.4  liters.  If  a  coefficient  appears  in  front  of  the  formula  of  a 
gas,  it  represents  the  number  of  gram-molecular  volumes.  Thus  2CO2 
stands  for  2V  (44.8  liters)  ;  2NH3  also  stands  for  2V.  How  many  liters 
of  gas  would  be  represented  by  4CO,?  by 



In  these  problems,  the  weight  of  one  substance  is  given  and 
the  volume  of  another  is  to  be  found.  Or  the  volume  of  one 
is  given,  and  the  weight  of  another  is  to  be  found.  The  pro- 
cedure is  the  same  in  both  cases.  Standard  conditions  of  tem- 
perature and  pressure  (S.T.P.)  are  assumed. 

EXAMPLE:  How  many  liters  of  nitric  oxide  can  be  prepared 
by  action  of  sufficient  dilute  nitric  acid  on  127.2  g.  of  copper? 


1.  Write  the  balanced  equation. 

3Cu  +  8HNO3  -»  3Cu(NO3)2  4-  4H2O  -f  2NO 

2.  Write  the  given  weight  over  its  formula  and  x  1.  over  the 
formula  whose  volume  is  to  be  found.  Cross  out  all  other 

127.2  g.  x  1. 

3Cu  +-SHN03  ->4€«fN0i);  +-4HsO-+  2NO 


3.  Under  the  formula  whose  weight  is  given,  write  its  molec- 
ular weight.  Under  the  formula  whose  volume  is  to  be  found, 
write  its  gram-molecular  volume  (V) ,  not  its  molecular  weight. 

127.2  g.  x  1. 

3Cu  -»  2NO 

3(63.6)  -  190.8       2V  -  2(22.4)  =  44.8 

4.  Write  the  mathematical  equation  based  on  the  relation- 

wt.  of  substance  used  vol.  of  substance  formed 

mol.  wt.  of  substance  used        V  of  substance  formed 

Solve  for  x.  127.2  _   xl. 

190.8  "  44.8 
190.8*  =  127.2(44.8) 

x  =  29.9  liters  of  NO 

The  same  general  procedure  is  followed  in  finding  the 
weight  of  one  substance  when  the  volume  of  another  is  given 
except  that  x  represents  the  unknown  weight  rather  than  the 
unknown  volume.  Use  an  equation  based  on  the  same  rela- 
tionship for  your  final  solution. 


1.  What  volume  of  H2  may  be  obtained  by  the  electrolysis 
of  90  g.  of  H2O? 

2.  How  many  liters  of  NH3  can  be  formed  by  the  action  of 
33  g.  of  (NHJ  2SO4  on  sufficient  Ca  (OH)  2? 

3.  How  much  NaCl  is  needed  to  produce  112  1.  of  HC1  gas? 

Nad  +  H2S04  ->  NaHSO4  +  HC1 

4.  What  weight  of  H2O  must  be  decomposed  to  produce 
224  ml.  of  O2? 

5.  A  manufacturer  requires  10,000  1.  of  N2O.  What  weight 
of  NH4NO3  must  be  decomposed? 


This  type  of  problem  involves  finding  the  volume  of  one  gas 
or  vapor  when  jhe  volume  of  another  is  known.  As  we  have 
learned,  the  coefficient  before  the  formula  of  a  gas  represents 
the  number  of  gram-molecular  volumes  of  the  gas.  Since  we 
are  dealing  only  with  volumes,  weights  are  disregarded.  Only 
the  volumes  as  represented  by  the  coefficients  are  considered. 



EXAMPLE:  How  many  liters  of  carbon  dioxide  are  formed  dur- 
ing complete  combustion  of  seven  liters  of  benzene,  C  H  ? 

'         66 


1.  Write  the  balanced  equation. 

2C6H6  +  15O2  -»  12CO2  +  6H2O 

2.  Write  the  given  volume  above  its  formula.  Write  x  L 
above  the  formula  whose  volume  is  to  be  found.  Cross  out  all 
other  formulas.  ,   " 

7  liters  x  I. 

3.  Write  the  number  of  gram-molecular  volumes  (shown  by 
the  coefficients)   under  the  respective  formulas. 

7  liters         x  1. 
2C6H6  -»  12C02 

2  12 

4.  Write  out  the  mathematical  equation  based  on  the  re- 

_  vol.  of  substance  used  vol.  of  substance  formed 

coefficient  of  substance  used      coefficient  of  substance  formed 

Solve  for  x.  7  _  x  1. 

2  "12" 
2*  =  84 
x  -  42  liters  of  CO2 


1.  50  1.  of  H2  react  completely  with  C12.  What  volume  of 
HC1  gas  is  formed? 

2.  What  volume  of  H2  is  necessary  to  unite  with  5  1.  of  O2 
without  leaving  any  O2  in  excess? 

3.  What  volume  of  NH8  can  be  made  from  5000  1.  of  pure 

4.  How  many  liters  of  O2  will  be  used  during  the  complete 
combustion  of  500  ml.  of  methane,  CH4? 

GH4  +  202  ->  G02  +  2H20 

5.  What  volume  of  oxygen  will  convert  50  ml.  of  NO  into 
nitrogen  dioxide,  NO2? 



Jaffe,  Bernard.  Chemical  Calculations.  World  Book  Co., 
Yonkers,  N.Y.,  1947.  Additional  problems  of  the  types  dis- 
cussed in  this  chapter,  together  with  a  more  detailed  account 
of  methods  of  determining  atomic  weights  and  molecular 

Jaffe,  Bernard.  Crucibles:  The  Story  of  Chemistry f  pp.  157- 
174.  Simon  and  Schuster,  New  York,  1948.  "The  Spirit  of  a 
Dead  Man  Leads  a  Battle"  tells  the  story  of  Avogadro. 

Perrin,  Jean  B.  Atoms.  D.  Van  Nostrand  Co.,  New  York, 
1923.  In  this  book,  Perrin,  who  won  the  Nobel  prize  for  his 
work  on  the  molecule,  describes  his  verification  of  Avogadro's 


1.  A  molecule  is  the  smallest  part  of  either  an  element  or  a 
compound  which  has  the  properties  of  that  substance. 

2.  Gay-Lussac's   law  states  that   the  relation   between   the 
combining  volumes  of  gases  and  the  volumes  of  their  products, 
if  gaseous,  may  be  expressed  in  small  whole  numbers. 

3.  Boyle's  law  states  that  the  volume  of  a  gas  varies  in- 
versely as  the  pressure  exerted  upon  it  if  the  temperature  re- 
mains constant. 

V1/V2  =  P2/P! 

4.  Charles'  law  states  that  the  volume  of  a  gas  varies  directly 
as  the  absolute  temperature  if  the  pressure  on  the  gas  remains 

V1/V2  =  Ti/T, 

5.  Avogadro's  law  states  that  equal  volumes  of  all  gases 
under  the  same  conditions  of  temperature  and  pressure  are 
composed  of  the  same  number  of  molecules. 

6.  Avogadro's  law  is  valuable  because   (1)  it  shows  that  the 
molecules  of  certain  elementary  gases,  among  them  hydrogen, 
oxygen,  nitrogen,  and  chlorine,  consist  of  two  atoms;    (2)  it 
makes  possible  the  determination  of  the  molecular  weights  of 
gases;    (3)  it  makes  possible  the  determination  of  the  atomic 
weights  of  gaseous  elements;  and    (4)  it  shows  the  relation- 
ships among  several  apparently  conflicting  facts  concerning  the 
actions  of  gases. 

7.  Avogadro's  hypothesis  was  verified  by  the  work  of  other 
scientists  and  today  is  a  chemical  law.  The  actual  number  of 


molecules  in  the  gram-molecular  weight  of  a  gas  was  deter- 
mined by  experiment.  This  number,  called  Avogadro's  number, 
is  the  same  for  all  gases.  It  is  6.02  X  1Q28- 

8.  The  gram-molecular  volume  of  a  gas  is  the  volume  occu- 
pied by  its  gram-molecular  weight.  Under  standard  conditions, 
it  is  22.4  liters. 


Group  A 

1.  In  1860,  what  was  the  condition  of  chemical  usage  with 
respect  to  the  terms  atom  and  molecule? 

2.  What  is  the  difference  between  an  atom  and  a  molecule? 

3.  When  and  by  whom  was  the  term  molecule  first  clearly 

4.  What  did  Cannizzaro  do  to  establish  the  meaning  of 

5.  (a)  State  Gay-Lussac's  law  and  (b)  give  two  illustrations 

ofit'  ' 


6.  The  volume  of  a  gas  changes  from  10  to  5  1.  when  the 
pressure  on  it  changes  from  1  to  2  atm.  What  law  does  this 

7.  The  volume  of  a  gas  changes  from  4  1.  to  2  1.  when  its 
temperature  changes  from  500°A.  to  250°A.  State  the  law 

8.  State  Avogadro's  law. 

9.  Upon  what  facts  did  Avogadro  base  his  hypothesis? 
10.  State  two  ways  in  which  Avogadro's  law  is  valuable. 

1 1 .  What  two  scientists  verified  Avogadro's  hypothesis? 

12.  What  is  a  gram-molecular  volume  of  a  gas? 

13.  Outline  the  method  used  in  working  a  weight- volume 

14.  Outline  the  method  used  in  working  a  straight- volume 

15.  If  15  1.  of  N2  are  needed  to  unite  with  O2  in  forming 
NO,  what  volume  of  O2  will  be  used? 

t  .  .  . 


16.  Assume  that  air  contains  20  percent  O2  by  volume.  What 
volume  of  air  will  be  needed  in  forming  100  ml.  of  O3? 

17.  What  volume  of  air  will  be  needed  for  the  complete 
combustion  of  750  ml.  of  acetylene,  C2H2? 

18.  CO  passed  over  warm  Ca  (OH)  2  reacts  as  follows: 

CO  +  Ca(OH)2  -»  CaCO3  +  H2  | 

How  does  the  volume  of  CO  compare  with  that  of  the  H2? 

19.  What  weight  of  carbon  is  in  44.8  1.  of  CO? 

20.  What  volume  of  NO2  will  be  formed  by  the  complete 
reaction  of  100.5  g.  of  Hg  with  concentrated  HNO3? 

Hg  +  4HN03  ->  Hg(N03)2  +  2H2O  +  2NO2 

21.  HC1  gas  was  bubbled  through  a  solution  of  NaOH.  As 
a  result,  468  g.  of  NaCl  were  formed.  What  volume  of  the 
HC1  used  actually  combined  with  the  base? 

Group  B 

22.  From  the  experimental  fact  that  3  vol.  of  O2  change  into 
2  vol.  of  O3  when  an  electric  discharge  is  passed  through  moist 
O2,  prove  that  the  molecule  of  ozone  contains  three  atoms. 

23.  Prove  nitrogen  molecule  contains  at  least  two  atoms. 

24.  Which  is  more  economical  to  use  in  the  preparation  of 
NH3,  (NH4)  2SO?  at  $8.75  per  100  Ib.  or  NH^Cl  at  12^  per  lb.? 

25.  (a)  Explain  the  operation  of  the  atomic-hydrogen  torch, 
(b)  Account  for  the  extreme  heat  obtained. 

26.  What  is  meant  by  absolute  temperature? 

27.  Prove  by  calculation  that  the  ounce-molecular-volume 
of  any  gas  equals  22.4  cu.  ft. 


1.  In  England,  Robert  Boyle  is  considered  the  father  of 
modern  chemistry.  In  France,  Lavoisier  is  called  the  father  of 
modern  chemistry.  Can  you  suggest  reasons  for  this  difference 
of  opinion  among  French  and  English  scientists?  Is  it  justified? 
Does  it  illustrate  scientific  open-mindedness? 

2.  Suggest  possible  reasons  for  the  neglect  of  Avogadro's 
hypothesis  from  1811   to  1860.  Can  you  cite  other  scientific 
work  which  remained  unrecognized  for  a  long  time? 

3.  Construct  a  cardboard  or  wooden  box  to  represent  the 
gram-molecular  volume  of  any  gas  at  S.T.P.  Exhibit  it  to  your 
class  alongside  of  a  quart  bottle  or  carton  of  milk. 





.  .  .  Sulfur  has  been  taken  intermit- 
tently from  Popocatepetl  since  the 
time  of  the  ancient  Aztecs,  who  used 
it  for  medicinal  purposes.  Two  of 
Cortez's  soldiers,  in  the  sixteenth 
century,  climbed  to  the  crater  and 
obtained  sulfur  for  the  purpose  of 
manufacturing  gunpowder.  Science 
News  Letter 

An  American  pharmacist  creates  a  new  industry.  The  discovery 
of  petroleum  in  Pennsylvania  in  1859  led  at  once  to  a  wide  search 
for  other  stores  of  oil.  Only  six  years  later,  oil  prospectors  stumbled 
upon  huge  deposits  of  sulfur,  a  yellow,  brittle,  lustrous  solid  known 
since  ancient  times.  These  deposits  were  about  500  feet  below  the 
surface  in  Louisiana  not  far  from  the  Gulf  near  the  Texas  line.  Here 
ages  ago  a  vast  geyser  had  spouted,  leaving  the  sulfur  within  and 
about  its  crater.  The  sulfur  was  covered  with  strata  of  clay,  limestone, 
and,  worst  of  all,  gas  and  quicksand.  It  was  impossible  to  sink  shafts 
to  reach  the  deposits  in  order  to  dig  the  sulfur  out  as  coal  is  mined. 
Many  companies  were  formed  to  exploit  these  deposits,  but  because 
of  the  many  difficulties,  they  all  failed. 

In  1891  Herman  Frasch  (frash)  heard  about  this  sulfur.  He  had 
come  from  Germany  25  years  before.  Leaving  high  school  at  16,  he 
had  been  apprenticed  to  a  druggist  and  then  had  left  for  the  United 
States,  where  he  opened  his  own  drugstore  in  Philadelphia.  Chem- 
istry fascinated  him,  and  in  the  back  of  his  drugstore  he  carried  on 
many  researches  on  petroleum  products.  Later  he  sold  his  store  and 
devoted  all  his  time  to  chemical  engineering.  He  tackled  the  problem 
of  extracting  sulfur  from  the  Louisiana  deposits. 



u      alr 


melted  sulfur 

'  '•'  :•  /'  :' V"'.\ :  '••'.'•"  anhydrite  •  "•  / ,'v 

Fig.  64.  The  Frasch  process  for  extracting  sulfur  from  deposits  deep  in  the  earth. 

When  Frasch  described  to  some  of  his  friends  the  entirely  new 
process  that  he  had  devised,  they  thought  it  impossible.  One  promi- 
nent person  challenged  him  in  no  uncertain  terms.  He  said  that  he 
would  swallow  every  ounce  of  sulfur  Frasch  extracted  by  his  queer 
process  of  pumping  a  solid  out  of  the  earth.  But  Frasch  kept  on  im- 
proving his  method  and  ultimately  he  succeeded  in  founding  a  new 
American  industry  based  on  his  process. 

The  Frasch  process  of  extracting  sulfur.  Frasch' s  plan  was  to  sink 
a  well  by  means  of  an  oil-drilling  rig,  and  lower  three  concentric 
pipes  (inside  a  casing)  down  to  the  sulfur.  Through  the  outermost 
6-inch  pipe,  superheated  water  was  to  be  pumped  to  melt  the  sulfur. 
Through  the  innermost  one-inch  pipe,  compressed  air  was  to  be 
forced  down  to  the  sulfur  deposit  to  make  the  melted  sulfur  frothy 
and  light.  The  result  was  to  be  a  flood  of  molten,  foamy  sulfur  gush- 
ing under  pressure  from  the  three-inch  pipe  between  the  other  two. 

Frasch  was  visibly  nervous  when  he  gave  the  order  to  start  the 
first  pump.  More  and  more  slowly  went  the  engine  with  its  increas- 
ing load  until  the  man  at  the  throttle  sang  out  at  the  top  of  his  voice, 
"She's  a-pumping."  A  liquid  appeared  at  the  mouth  of  the  three-inch 
pipe.  Frasch  wiped  the  liquid  off  a  polished  piston  rod,  and  gazed 
upon  the  first  crystals  of  sulfur  extracted  by  his  ingenious  process. 
Then  came  a  steady  stream  of  golden  liquid,  which,  in  15  minutes, 
filled  every  one  of  the  40  barrels  available.  Still  the  molten  sulfur 
kept  pouring.  Embankments  were  quickly  thrown  up  and  lined  with 
boards  to  hold  the  sulfur  as  it  solidified. 

This  is  the  method  of  mining  sulfur  in  use  today  in  Texas  and 
Louisiana.  These  deposits  now  supply  much  of  the  world's  needs.  In 
this  way,  mountains  of  sulfur,  99^  percent  pure,  are  extracted  and 
stand  ready  to  be  dynamited  into  pieces  small  enough  to  be  loaded 
and  shipped  to  many  parts  of  the  world. 



Sulfur  is  obtained  in  Sicily  by  a  different  process.  Before  the  time 
of  Frasch,  a  group  of  English  financiers  had  been  marketing  the  rich 
sulfur  deposits  of  the  volcanic  region  of  Sicily.  Here  the  sulfur  is 
found  mixed  with  clay  and  limestone,  from  which  it  is  separated  by 
melting  this  ore,  allowing  the  free  sulfur  to  flow  away  from  the  im- 
purities. The  sulfur  is  then  boiled  and  is  changed  to  a  powder,  called 
(loiuers  of  sulfur,  by  chilling  the  vapors. 

The  English  monopoly  had  been  supplying  more  than  90  percent 
of  the  world's  sulfur.  In  1904,  when  American  output  reached  the 
point  where  a  single  well  could  supply  400  tons  of  sulfur  daily,  the 
English  company  went  out  of  business.  To  prevent  the  unemploy- 
ment of  hundreds  of  workers,  the  Italian  government  took  control. 
Frasch  aided  gladly  in  stabili/ing  the  Sicilian  sulfur  industry. 

Crude  free  sulfur  accounts  for  about  only  40  percent  of  the  world's 
sulfur  production.  Iron,  zinc,  and  copper  sulficles,  natural  gas,  and 
industrial  gases  supply  the  rest.  The  United  States  today  leads  the 
world  in  the  production  of  sulfur.  Many  countries,  including  Italy 
and  Mexico,  are  stimulating  their  own  production  of  sulfur  by  sub- 
sidies, tariff  laws,  bounties,  and  price  guarantees. 

Physical  properties  of  sulfur.  Sulfur  is  a  pale  yellow,  soft,  brittle 
solid  (plastic,  or  pliable,  in  the  case  of  amorphous  sulfur)  with  a 
very  faint  odor  and  no  marked  taste.  It  is  practically  insoluble  in 
water,  more  soluble  in  carbon  tetrachloride,  CC14,  and  very  soluble 
in  carbon  disulfide,  CS... 

Sulfur  is  a  poor  conductor  of  heat.  It  melts  at  about  114°C., 
forming  a  pale  yellow  liquid,  which  on  further  heating  darkens  and 
thickens,  becoming-  almost  black  at  235 °C.  At  a  still  higher  tempera- 
ture, it  becomes  thinner  again,  and  finally  changes  to  a  yellow  vapor 
at  445°C.  Sulfur  is  mentioned  in  the  Bible  as  brimstone. 

Texas  Gulf  Sulfur  Cohipany 

Liquid  sulfur  flows  into  the 
storage  vat  in  which  it  will 
solidify.  The  pipeline  leads 
to  the  sulfur  well. 



Allotropic  forms  of  sulfur.  Oxygen,  as  you  have  learned,  occurs  in 
two  allotropic  forms,  O2  and  O3,  but  sulfur  occurs  in  two  common 
crystalline  forms  (rhombic  and  prismatic) ,  and  one  noncrystalline 
or  amorphous  form.  Each  has  different  properties  caused  by  differ- 
ences in  the  arrangement  of  their  atoms. 

1)  Rhombic  sulfur.  Sulfur  is  found  in  nature  in  the  form  of 
rhombic  crystals,  the  most  stable  form  under  normal  conditions. 
Their  molecules  consist  of  rings  of  eight  atoms  of  sulfur   (Ss) ,  and 
their  density  is  about  two.  Crystals  of  this  form  of  sulfur  may  be 
prepared  by  dissolving  sulfur  in  carbon  disulftde  and  allowing  the 
solvent  to  vaporize  slowly.  The  residue  consists  of  perfect  crystals 
having  the  shape  of  two  pyramids  with  their  bases  joined.  Roll  sulfur, 
made   by  cooling  sulfur   in  cylindrical   molds,   is   almost   entirely 

2)  Prismatic  sulfur.  When  sulhir  is  heated  until  it  just  melts  and 
is  then  allowed  to  cool  slowly,  it  forms  long  needle-shaped  crystals 
whose  density  is  somewhat  less  than  that  of  rhombic  sulfur. 

3)  Amorphous  sulfur.  When  boiling  sulfur  is  suddenly  cooled  by 
pouring  it  into  cold  water,  the  rings  are  broken  and  are  replaced  by 
long  chains  of  sulfur  atoms  which  become  entangled  and  can  be 
stretched  like  rubber.  It  is  amber  in  color  and,  unlike  the  other 
two    forms,    is    insoluble    in    carbon    disulfide.    Amorphous    sulfur 
changes  gradually  into  rhombic  at  room  temperature.  For  simplicity, 
all  forms  are  designated  S.  Flowers  of  sulfur  is  a  powdered  mixture 
composed  of  rhombic  and  some  of  the  plastic.  Between  96 °C.  and 
1 14°C.  the  most  stable  form  of  sulfur  is  the  prismatic. 

Chemical  properties  of  sulfur.  The  atomic  weight  of  sulfur  is 
32.  Its  atomic  number  is  16,  and  it  has,  therefore,  six  electrons  in  its 
third,  or  outermost,  ring.  This  makes  it  a  borrower  of  electrons.  It 
is  a  nonmetal,  fairly  active,  and  has,  under  ordinary  conditions,  a 
valence  of  minus  two.  Therefore,  sulfur  closely  resembles  oxygen  in 
its  chemical  properties. 

Fig.  65.  Allotropic  form,  of  sulfur.  melfed  Sulfur" 



Exterior  of  a  sulfur  storage  vat. 
The  solidified  sulfur  will  be 
blasted  into  small  pieces  for 
shipment  in  railroad  cars. 

Texas  Gulf  Sulfur  Company 

Like  oxygen,  sulfur  unites  with  most  metals.  The  compounds  of 
metals  and  sulfur  are  called  sulfides.  For  example,  when  sulfur  is 
heated  with  iron  powder,  iron  sulfide  is  formed. 

-        -   •  Fe-fS-^FeS 

Hot  copper  burns  in  sulfur  vapor,  forming  cuprous  sulfide. 

.    /-   V^    •    •  2Cu+S-+Cu2S 

When  sulfur  is  mixed  with  zinc  dust  and  ignited,  the  chemical 
reaction  is  so  vigorous  that  a  great  deal  of  light  and  heat  are  liber- 
ated and  dense  clouds  of  zinc  sulfide,  which  settle  out  as  a  powder, 
are  formed. 

Zn  +  S-^  ZnS 

Although  sulfur  is  a  nonmetal,  it  is  less  nonmetallic  than  oxygen. 
It  therefore  combines  with  oxygen,  exhibiting  a  positive  valence  of 
either  four  or  six.  Sulfur  burns  in  air  with  a  pale  blue  flame,  form- 
ing sulfur  dioxide,  SO.,.  A  small  amount  of  sulfur  trioxide,  SO3,  is 
formed  later  by  further  oxidation  of  the  SO2. 

2SO2  +  O2 

>  SO2  (sulfur  dioxide) 

>  2SO3  (sulfur  trioxide) 

The  electron  structure  of  sulfur  and  oxygen  are  shown  below. 
Fig.  66.  The  structure  of  the  oxygen  atom  (left)  and  the  sulfur  atom. 





Sulfur  also  shows  more  metallic  properties  than  chlorine  by  unit- 
ing with  the  latter,  forming  sulfur  dichloride,  SCI.,,  a  brownish  red 
liquid  with  a  pungent  odor  used  in  chlorinating,  and  also  sulfur 
monochloride,  a  heavy,  amber-colored  unstable  liquid  with  an  irri- 
tating odor,  used  in  vulcanizing  rubber. 

2S  +  C12  ->  S2C12  (sulfur  monochloride) 

When  sulfur  vapor  is  passed  over  carbon  heated  in  an  electric  fur- 
nace, the  two  elements  combine,  forming  carbon  disulfide,  CS2. 

C  +  2S  -»  CS2 

Carbon  disulfide  is  a  heavy,  colorless  liquid  with  a  pleasant  odor. 
Generally,  however,  it  has  a  disagreeable  odor  caused  by  impurities. 
It  is  very  combustible.  Its  chief  use  is  as  a  solvent  for  sulfur,  gums, 
rubber,  fats,  and  waxes.  It  has  been  used  also  as  a  poison  in  exter- 
minating ants  and  other  insects,  rats,  and  prairie  dogs. 

Rubber  tires  depend  on  sulfur.  The  normal  annual  consumption 
of  sulfur  in  the  United  States  is  more  than  three  million  tons.  It  is 
one  of  the  fundamental  industrial  elements.  By  far  the  greatest 
quantities  of  sulfur  are  used  in  the  manufacture  of  sulfur ic  acid, 
one  of  the  most  widely  used  of  the  heavy  chemicals.  The  production 
and  uses  of  sulfuric  acid  are  discussed  in  Chapter  21. 

Great  quantities  of  sulfur  are  used  in  vulcanizing  rubber.  By  this 
process,  soft,  sticky,  perishable,  natural  rubber  is  changed  to  a  harder, 
less  plastic,  more  durable  rubber.  Vulcani/ed  rubber  is  used  chiefly 
in  making  automobile  tires  but  also  in  making  thousands  of  other 
rubber  articles. 

Tfie  Firestone  Tire  and  Rubber  Company 

In  a  modern  tire  factory,  a  worlc- 
man    removes    a    finished    tire 
•om    a    steam-heated    vulcani- 
zing mold. 


When  the  American,  Charles  Goodyear,  in  1839  while  working  in 
his  kitchen  in  Woburn,  Massachusetts,  accidentally  dropped  a  piece 
of  rubber  mixed  with  sulfur  on  a  hot  stove,  he  discovered  the  proc- 
ess of  vulcanization  (derived  from  Vulcan,  the  Roman  god  of  fire) . 
A  new  and  highly  important  industry  was  made  possible.  To  shorten 
the  time  required  for  vulcanizing  rubber,  a  catalyst,  or  accelerator, 
such  as  zinc  oxide,  is  added  to  the  mixture  of  rubber  and  sulfur. 

In  making  an  article  of  rubber,  the  washed  raw  rubber  is  thor- 
oughly mixed  with  various  chemicals  that  determine  the  properties 
of  the  finished  product.  Among  these  substances  are  sulfur,  carbon 
black,  lead  oxide  (PbO) ,  zinc  oxide,  and  carbonates  of  magnesium 
and  calcium.  Different  combinations  of  these  and  other  substances 
in  varying  quantities  may  be  used  in  accordance  with  the  properties 
desired  in  the  finished  product.  The  rubber  is  then  rolled  into  sheets 
of  the  desired  thickness  or  placed  in  molds  of  the  desired  shape. 
While  in  the  mold,  the  rubber  is  heated  and  vulcanization  occurs. 

Sulfur  is  used  in  controlling  fungi  and  insects.  A  very  effective 
liquid  for  destroying  fungus  growths  and  fruit  tree,  shrub,  and  vine 
pests  is  a  deep  orange-red  lime-sulfur  spray  made  by  boiling  sulfur 
with  calcium  hydroxide.  The  spray  is  both  a  fungicide,  a  substance 
that  kills  fungi  and  molds,  and  an  insecticide,  a  substance  that  kills 
insects.  Dusting  with  very  finely  powdered  sulfur  is  effective  against 
rose  diseases,  mildew,  and  black  spot.  However,  ordinary  flowers  of 
sulfur  is  not  fine  enough  to  be  of  much  value  as  a  dusting  agent,  and 
even  some  finely  ground  commercial  dusting  powders  are  too  coarse. 

Colloidal  sulfur  and  wettable  sulfur,  the  first  a  very  highly  dis- 
persed sulfur  in  water  and  the  second  sulfur  so  treated  that  it  dis- 
perses on  contact  with  water,  are  both  used  in  making  mild  sprays 
that  are  particularly  useful  for  the  summer  spraying  of  roses  and  for 
the  control  of  mildew  and  true  rust  diseases  of  other  plants. 

Sulfur  is  used  in  medicine.  Ointments  containing  sulfur  have 
been  used  since  antiquity  to  control  skin  diseases  caused  by  fungi. 
A  common  ointment  of  this  kind  is  made  by  mixing  three  parts  by 
weight  of  sulfur,  15  parts  of  white  petrolatum,  or  Vaseline,  one  part 
of  lanolin,  or  wool  fat,  and  one  part  of  yellow  wax.  Such  an  oint- 
ment is  effective  in  killing  the  very  small  mites  that  cause  scabies, 
or  itch.  Sulfur  ointments  may  be  of  value  in  the  control  of  infectious 

Other  uses  of  sulfur.  Much  sulfur  goes  into  the  manufacture  of 
calcium  bisulfite,  Ca  (HSO3)  2,  which  is  used  in  making  wood  pulp 
for  the  manufacture  of  paper.  Sulfur  is  used  also  in  making  syn- 
thetic resins,  sulfur  colors,  and  gunpowder.  Sulfur-lined  steel  pipes 



are  used  to  transport  liquids  that  are  very  corrosive  to  the  materials 
of  which  pipes  are  normally  made.  Sulfur  cements  are  used  to  join 
bricks  in  floors  and  walls  that  are  continually  subjected  to  the  corro- 
sive effects  of  acids  or  alkalies.  Great  quantities  of  sulfur  are  used 
in  the  making  of  matches. 

How  matches  are  made.  Early  matches  were  dangerous,  hard  to 
use,  and  hard  to  carry.  Phosphorus  matches  often  caused  painful 
burns  and  were  harmful  to  the  health  of  the  workers  who  made  them. 

The  first  friction  match  was  invented  in  1827  by  an  English  phar- 
macist, John  Walker.  It  contained  potassium  chlorate  (KC1O3) ,  an 
oxidi/ing  agent,  and  antimony  sulfide  (Sb,Ss) ,  a  compound  with  a 
low  kindling  temperature.  The  locofoco  match  was  an  American 
adaptation  of  Walker's  friction  match.  During  the  presidential  cam- 
paign of  1840,  the  Democrats  were  called  locofocos  because  at  one  of 
their  meetings,  they  used  matchlight  when  the  Whig  landlord  turned 
off  the  gas. 

In  1831  white  phosphorus  was  used  for  the  first  time  in  making 
matches.  It  was  a  more  efficient  fire-producer  than  antimony  sulfide, 
but  played  havoc  with  the  health  of  the  match-factory  workers,  who 
finally  demanded  protection.  This  led  in  1906  to  an  international 
convention  which  prohibited  the  further  use  of  white  phosphorus 
in  making  matches  because  of  its  poisonous  nature. 

The  strike-anywhere  match  in  use  today  also  contains  both  an  oxi- 
dizing agent  and  a  compound  with  a  low  kindling  temperature.  The 
head  of  the  paraffin-dipped  matchstick  contains  potassium  chlorate 
and  phosphorus  sulfide  (P4S3) ,  a  dark  solid  that  is  concentrated  on 
the  tip  of  the  match.  Glue  is  used  to  bind  the  chemicals,  and  ground 
glass  or  other  abrasive  is  sometimes  added  as  a  filler. 

The  tip  of  a  safety  match  is  composed  of  easily  combustible  anti- 
mony sulfide  (Sb.,S3)  ,  and  potassium  chlorate,  which  provides  addi- 
tional oxygen.  The  side  of  the  box  contains  red  phosphorus,  which 



hydrogen  sulfide  S 

Fig.  67.  Laboratory  preparation  of 
hydrogen  sulflde  (left).  Structure 
of  the  hydrogen  sulflde  molecule 

is  nonpoisonous.  The  material  on  the  tip  of  the  match  will  not  ig- 
nite easily  unless  it  is  rubbed  on  the  treated  side  of  the  box. 

Disastrous  fires  are  frequently  caused  by  matches  that  have  been 
blown  out  but  still  retain  glowing  tips.  To  prevent  this  dangerous 
afterglow,  matches  are  dipped  in  some  such  solution  as  sodium  sul- 
late  or  ammonium  phosphate  during  their  manufacture.  When 
blown  out,  they  do  not  leave  glowing  tips. 

Waterproof  matches  are  treated  with  a  transparent  coating  with 
a  high-kindling  temperature.  They  will  light  even  after  being  soaked 
in  water. 

The  manufacturer  of  matches  uses  a  continuous-process  machine, 
which  takes  pine  wood,  cuts  it,  dips  the  sticks  in  paraffin,  coats  them 
with  the  chemicals  needed,  dries  them,  and  finally  packs  them  for 
shipment.  More  than  a  million  matches  can  be  made  by  one  of 
these  machines  each  hour. 

How  sulfur  and  hydrogen  unite.  Hydrogen  has  only  one  electron. 
It  therefore  needs  one  electron  to  complete  its  only  ring;  but  since 
it  shows  a  strong  tendency  to  lend  its  one  electron,  it  is  said  to  pos- 
sess metallic  properties.  Sulfur,  requiring  two  electrons  to  complete 
its  outer  ring,  will  combine  with  two  atoms  of  hydrogen.  This  union 
forms  the  compound,  hydrogen  sulfide,  H,S,  as  shown  in  Fig.  67. 
In  this  compound,  sulfur  exhibits  definitely  nonmetallic  properties. 
Although  it  is  possible  to  prepare  hydrogen  sulfide  by  direct  union 
of  sulfur  and  hydrogen,  it  is  prepared  most  easily  by  other  methods. 


How  hydrogen  sulfide  is  prepared.  Hydrogen  sulfide  is  prepared 
both  commercially  and  in  the  laboratory  by  the  general  method  of 
preparing  an  acid,  that  is,  by  the  action  of  sulfuric  acid  on  a  salt  (sul- 
fide) .  The  sulfide  most  generally  used  is  ferrous  sulfide,  FeS,  a  black 
iron  ore.  When  sulfuric  acid  is  poured  on  ferrous  sulfide  in  a  test 
tube  (see  Fig.  67) ,  hydrogen  sulfide  gas  is  given  off  immediately 
without  the  addition  of  any  external  heat. 

FeS  -I-  H2SO4  ->  FeSO4  +  H2S  | 

Physical  properties  of  hydrogen  sulfide  gas.  Hydrogen  sulfide  is 
colorless,  slightly  heavier  than  air,  and  fairly  soluble  in  water  (one 
volume  of  water  dissolves  three  volumes  of  hydrogen  sulfide)  .  Natu- 
ral sulfur  waters  contain  hydrogen  sulfide  in  solution  and,  upon  be- 
ing decomposed,  leave  a  deposit  of  free  sulfur.  Easily  liquefied, 
hydrogen  sulfide  is  sold  in  cylinders  for  laboratory  use.  Its  most  char- 
acteristic physical  property  is  its  odor,  the  odor  of  rotten  eggs.  In 
fact,  it  is  hydrogen  sulfide  that  gives  such  eggs  their  odor.  It  is  caused 
by  the  decomposition  of  organic  sulfur  compounds  in  the  yolks. 

Hydrogen  sulfide  forms  naturally  in  marshes,  oil  wells,  mines  and 
coal  piles,  manure  pits,  and  sewers.  In  persons  who  have  endured 
mild  exposure  to  its  effects,  it  produces  inflamed  thrpat,  headache, 
a  heavy  feeling  in  the  stomach,  and  dizziness.  When  breathed  in 
large  quantities  it  causes  death.  Both  natural  gas  and  coal  gas  con- 
tain H2S  which  is  removed  before  they  are  used  as  household  fuels. 

Chemical  properties  of  hydrogen  sulfide  gas.  When  burned  in 
sufficient  air,  hydrogen  sulfide  gives  off  a  pale  blue  flame,  and  water 
vapor  and  sulfur  dioxide  are  formed.  This  sulfur  dioxide  gives  burn- 
ing hydrogen  sulfide  its  irritating  odor.  The  equation  for  the  com- 
plete combustion  of  hydrogen  sulfide  is: 

2H2S  +  3O2  -»  2H2O  +  2S02  T 

When  hydrogen  sulfide  is  burned  in  a  small  amount  of  air  (incom- 
plete combustion) ,  water  is  formed  as  before,  but  free  sulfur  is  pro- 
duced instead  of  sulfur  dioxide.  This  sulfur  separates  out  as  a  yellow 
powder.  The  fact  that  it  does  so  probably  accounts  for  the  presence 
of  sulfur  around  volcanoes,  which  emit  hydrogen  sulfide  gas. 

2H2S  +  O2  -»  2H2O  +  2S  T 

Because  of  the  tendency  of  hydrogen  sulfide  to  unite  with  oxygen, 
it  is  a  fairly  good  reducing  agent.  When  hydrogen  sulfide  acts  as  a 
reducing  agent,  the  sulfur  lends  electrons  and  is  oxidized. 



In  a  water  solution,  hydrogen  sulfide  dissociates  to  some  extent. 
Such  a  solution  acts  as  a  weak,  unstable  acid,  sometimes  known  as 
hydrosulfuric  acid.,  H2S.  On  continued  boiling,  this  acid  liberates  hy- 
drogen sulfide,  leaving  pure  water. 

Hydrogen  sulfide  reacts  with  certain  metals  and  also  with  the  salts 
of  certain  metals,  forming  sulfides.  The  tarnishing  of  silverware  is 
caused  by  the  formation  of  a  brownish-black  sulfide  of  silver,  Ag2S. 
The  blackening  of  lead  paints  is  caused  by  the  formation  of  black 
lead  sulfide.  Lithopone,  a  white  paint  base  now  used  widely,  consists 
of  a  mixture  of  barium  sulfate  (RaSO4)  ,  and  zinc  sulfide  (ZnS)  . 
Lithopoiie  paints  do  not  lose  their  color  by  the  action  of  H,S.  Why? 
Titanium  dioxide,  TiO2,  also  is  used  as  a  white  paint  base.  It  is  not 
blackened  by  sulfur  compounds  and  has  great  covering  power. 

Many  important  sulfides  are  found  in  the  earth.  The  salts  of 
hydrosulfuric  acid  form  an  important  class  of  compounds  called 
sulfides.  Many  of  them  occur  in  nature  and  constitute  important  ores 
such  as  iron  pyrites.,  FeS.,;  galena,  PbS;  zinc  blende,  ZnS;  cinnabar, 
HgS;  and  CuS.  Certain  colored  sulfides  are  used  as  mineral  pig- 
ments in  the  coloring  of  paints.  Cadmium  sulfide,  CdS,  for  example, 
is  a  yellow  pigment,  and  zinc  sulfide,  ZnS,  is  a  white  pigment  used 
in  paints  (see  illustration  following  page  382). 

The  chief  use  of  hydrogen  sulfide.  In  the  analysis  of  ores  and  in 
the  separation  of  groups  of  certain  metals  from  other  groups  of  met- 
als, hydrogen  sulfide  gas  is  indispensable.  For  this  reason,  a  hydro- 
gen sulfide  generator  is  always  present  in  a  laboratory  for  analytical 
chemistry.  The  sulfides  of  certain  metals  such  as  sodium  and  calcium 
are  soluble  in  water,  while  those  of  other  metals  such  as  lead  and  zinc 
are  insoluble. 

By  passing  hydrogen  sulfide  into  a  solution  of  the  soluble  salts  of 
such  metals,  the  sulfides  of  certain  metals  precipitate  out  and  may 
be  separated  by  filtration.  Thus,  certain  metals  present  in  an  ore  or 

Standard    Oil   Com  puny    (N  .J .) 

Hydrogen  sulfide  fumes  ris- 
ing from  crude  oil  may 
prove  fatal  to  workmen. 
Hence,  protective  masks 
are  worn  by  anyone  work- 
ing near  the  oil  storage 


alloy  may  be  separated  with  the  aid  of  H2S.  Furthermore,  since  the 
colors  of  the  sulfides  of  many  metals  differ,  chemists  can  use  these 
differences  in  color  in  detecting  the  presence  of  these  metals.  For 
example,  zinc  sulfide  is  white,  arsenic  sulfide  is  yellow,  antimony 
sulfide  is  orange,  and  copper  sulfide  is  brownish-black. 

Zn(NO3)2  +  H2S  ->  ZnS  I  +  2HNO3 
CuSO4  +  H2S  ->  CuS  I  +  H2SO4 

This  difference  in  color  is  only  one  of  the  very  many  ingenious 
methods  used  by  analytical  chemists  in  detecting  and  isolating  ele- 
ments present  in  complex  mixtures  and  compounds. 

The  test  for  the  sulfide  ion.  If,  on  the  addition  of  sulfuric  acid  to  a 
compound,  hydrogen  sulfide  gas  is  liberated,  the  substance  tested 
is  a  sulfide.  The  hydrogen  sulfide  liberated  is  easily  detected  either 
by  its  odor  or  by  its  ability  to  turn  a  silver  coin  brownish-black,  as  a 
result  of  the  formation  of  silver  sulfide. 


Fabre,  Jean  H.  The  Wonder  Book  of  Chemistry,  pp.  345-359. 
Albert  &T  Charles  Boni,  New  York,  1922.  Discusses  sulfur  and 
includes  some  simple  experiments  with  this  element. 

Waggaman,  W.  H.,  and  Barr,  J.  A.  "Sulfur  for  Survival." 
Chemistry,  October,  1951,  pp.  1-10.  An  illustrated  article  on 
sources,  extraction,  and  properties  of  sulfur. 


Group  A 

1.  When  and  by  whom  was  the  first  successful  method  of 
obtaining  S  from  the  Louisiana  deposits  invented? 

2.  What  was  the  great  difficulty  that  had  to  be  overcome 
before  S  could  be  extracted  from  the  Louisiana  deposits? 

3.  How  did  the  Frasch  process  affect  the  Sicilian  sulfur 

4.  (a)  Make  a  labeled  diagram  of  the  Frasch  process  and 
(b)  explain  it. 

5.  What  are  two  ways  in  which  the  Louisiana  sulfur  de- 
posits differ  from  the  Sicilian  sulfur  deposits? 

6.  Name  the  allotropic  forms  of  S  and  tell  how  each  may 
be  prepared  in  the  laboratory. 


7.  What  two  elements  other  than  S  occur  in  allotropic 

8.  (a)  State  the  chief  differences  and  resemblances  of  the 
three  allotropic  forms  of  S.  (b)  How  could  you  prove  that  all 
three  forms  of  S  are  the  same  element? 

9.  A  piece  of  plastic  S  is  left  overnight  in  the  laboratory. 
The  next  morning  yellow  brittle  S  is  found.  Explain. 

10.  (a)  How  are  the  two  kinds  of  commercial  sulfur,  roll 
sulfur  and  flowers  of  sulfur,  prepared?    (b)  What  allotropic 
forms  are  in  each? 

1 1 .  Make  a  diagram  of  an  atom  of  S  and  use  it  to  find  the 
valence  of  S,  to  describe  its  chemical  activity,  and  to  explain 
why  it  is  a  nonmetal. 

12.  In  what  three  ways  does  S  resemble  O2? 

13.  Write  balanced  equations  for  the:    (a)  union  of  sulfur 
and  zinc;  (b)  complete  combustion  of  hydrogen  sulfide;  (c)  in- 
complete combustion  of  hydrogen  sulfide;   (d)  burning  of  sul- 
fur; (e)  union  of  sulfur  and  chlorine;  (f)  complete  combustion 
of  carbon  disulfide. 

14.  What  are  four  chief  uses  of  S?  List  in  order  of  im- 

t .  . . 

15.  (a)  What    is    meant    by    the    vulcanizing   of    rubber? 

(b)  What    properties    does    vulcanizing    impart    to    rubber? 

(c)  What  American  discovered  this  process? 

16.  (a)  Of  what  chemicals  is  the  tip  of  a  strike-anywhere 
match  composed?  (b)  How  does  a  safety  match  differ  from  an 
ordinary  match? 

17.  Hydrogen  sulfide  present  in  natural  gas  is  removed  be- 
fore the  natural  gas  is  sent  into  the  pipelines  to  be  used  as  fuel. 

18.  By  a  labeled  diagram  describe  the  laboratory  prepara- 
tion of  H2S. 

19.  What  weight  of  FeS  would  be  needed  to  prepare  204  Ib. 
of  H2S? 

20.  What  volume  of  gaseous  H2S  could  be  prepared  from 
20  g.  of  90  percent  pure  FeS? 

21.  A  compound  of  hydrogen  and  sulfur  has  a  mol.  wt.  of  34. 
The  percentage  of  S  in  the  compound  is  94.1  percent.  Find  its 


22.  What  volume  of  H2S  is  required  to  precipitate  all  the 
CuS  from  a  solution  containing  80  g.  of  CuSO4? 

23.  What  volume  of  SO2  will  be  formed  by  the  complete 
combustion  of  896  ml.  of  gaseous  H2S? 

24.  Write  ionic  equations  for  the  following  and  tell  whether 
each  reaction  goes  to  completion  (refer  to  Table  10,  page  243) : 

a)  FeS  +  H2S04  -»  FeSO4  +  H2S 

b)  Pb(N03)2  +  H2S  -»  PbS  +  2HNO3 

c)  ZnSO4  +  (NH4)2S  ->  ZnS  +  (NH4)2SO4 

d)  Na2SO4  +  H2S  -»  Na2S  +  H2SO4 

25.  Compare  the  physical  properties  of  H2S  and  N2O. 

26.  Describe  what  happens  when  gaseous  H2S  is  bubbled 
through  water. 

27.  Why  is  H2S  said  to  be  (a)  weak?  (b)  unstable? 

28.  Explain  the  tarnishing  of  silverware. 

29.  Why  is  H2S  solution  kept  in  amber-colored  bottles? 

30.  Using  two  equations,  show  the  chief  use  of  H2S. 

31.  Describe  a  test  for  a  sulfide. 

32.  Copy  and  complete.  Do  not  write  in  this  book.  S  occurs 
in  our  bodies  because  it  is  one  of  the  elements  found  in  .... 
The  two  chief  sulfur-producing  states  in  America  are  Louisi- 
ana and  ....  A  lime-sulfur  preparation  is  used  as  ....  A  sulfur 
compound  used  in  exterminating  rats  is  ....  A  sulfide  that  is 
soluble  in  water  is  .... 

33.  Under  a  rubber  band  used  to  keep  pieces  of  silver 
cutlery  together,  a  black  mark  is  found.  Explain. 

Group  B 

34.  How  would  you  tell  fool's  gold,  FeS2,  from  genuine  gold? 

35.  In  the  Frasch  process,  why  is  S  forced  out  of  the  three- 
inch  pipe  rather  than  out  of  the  six-inch  pipe? 

36.  The  action  of  H2S  on  a  solution  of  ZnSO4  is  reversible. 
How  would  you  force  this  reaction  to  completion? 

37.  NH3  leaks  may  be  detected  by  burning  sulfur  candles. 

38.  (a)  Why  can  we  not  use  HNO8  in  preparing  H2S  from 
FeS?  (b)  What  is  formed  if  H2S  is  passed  into  HNO3? 


39.  Why  do  we  write  CS2  and  not  S2C? 

40.  A  mixture  of  sulfur  and  molasses  was  given  to  children 
as  a  "spring  tonic."  What  do  you  think  of  this  practice? 


1.  Discovery  of  the  vulcanization  of  rubber  is  generally  con- 
sidered an  accident.  Had  the  accident  not  happened  to  Good- 
year, would  the  thousands  of  rubber  articles  in  use  today  never 
have  come?  Explain  your  answer  fully  in  terms  of  scientific 
advance  and  the  needs  of  society. 

2.  If  you  live  in  a  rural  section,  talk  with  as  many  farmers 
as  possible  to  learn  what  methods  they  use  to  check  the  losses 
incurred   by  insect  pests   and  fungus  growths.   Report  your 
findings  to  your  class. 

3.  Filter  paper  dipped  in  a  solution  of  Pb  (C2H3O2)  2  turns 
black  in  the  presence  of  H2S.  Hang  small  strips  of  it  in  the 
basement  while  the  furnace  is  burning,  over  the  kitchen  stove 
or  gas  range  while  dinner  is  being  cooked,  and  in  the  living 
room.  Which  of  the  strips  turns  blackest  first?  What  do  you 




.  .  .  We  may  fairly  judge  of  the 
commercial  prosperity  of  a  country 
from  the  amount  of  sulfuric  acid  it 
consumes.  Reflecting  upon  the  im- 
portant influence  which  the  price  of 
sulfur  exercises  .  .  .  we  can  under- 
stand why  the  English  Government 
should  have  resolved  to  war  with 
Naples,  in  order  to  abolish  the  siu- 
fur  monopoly,  which  the  latter 
power  attempted  recently  to  estab- 
lish. Justus  yon  Liebig, 

Why  sulfuric  acid  is  the  fundamental  chemical  today.  One  of  the 

outstanding  differences  between  our  society  and  the  society  which 
preceded  it  is  the  tremendous  role  that  machines  play.  Automobiles, 
locomotives,  steel  ships,  and  airplanes  are  comparatively  new.  In  the 
manufacture,  operation,  and  upkeep  of  these  and  thousands  of  other 
machines,  sulfuric  acid  is  directly  or  indirectly  a  prime  necessity.  If 
the  ability  to  make  this  acid  were  lost  suddenly,  industry  would  be 
paralyzed.  So  fundamental  is  it  to  our  industrial  life  that  its  con- 
sumption is  a  fair  index  of  industrial  conditions. 

How  sulfur  dioxide  is  prepared  for  use  in  industry.  About  90  per- 
cent of  the  sulfur  consumed  industrially  is  first  burned  in  air  to 
produce  sulfur  dioxide.  Heating  sulfide  ores  is  another  commercial 
method  of  preparing  SO2.  When  an  ore,  such  as  iron  pyrites,  FeS2,  is 
heated  in  air,  or  roasted,  one  of  the  products  formed  is  sulfur  dioxide. 

4FeS2  +  11O2  -»  2Fe2O3  +  8SO2  T 

Sulfur  dioxide  occurs  in  small  quantities  in  the  vapors  rising  from 
active  volcanoes  and  certain  sulfur  Springs,  and  in  the  gases  formed 
during  the  burning  of  coal. 


Na2SO3  +  H2SO4 

Fig.  68.  Laboratory  prepa- 
ration of  sulfur  dioxide. 
Why  is  the  excess  gas 
passed  through  water? 

-*—  H2Q 

How  sulfur  dioxide  is  prepared  in  the  laboratory.  When  sodium 
sulfite,  Na2SO3,  or  any  other  sulfite  is  treated  with  an  acid,  sulfur 
dioxide  is  liberated.  The  gas  is  collected  by  the  displacement  of  air, 
as  shown  in  Fig.  68. 

Na2SO3  +  H2SO4  -»  Na2SO4  +  H2O  +  SO2 1 

Physical  properties  of  sulfur  dioxide.  Sulfur  dioxide  is  colorless, 
has  a  suffocating  odor,  is  more  than  twice  as  heavy  as  air,  and  is  very 
soluble  in  water  (one  volume  of  water  at  room  temperature  dissolves 
40  volumes  of  the  gas)  .  Under  a  pressure  of  only  three  atmospheres 
(approximately  45  pounds  per  square  inch,  or  45  psi)  it  changes  to 
a  colorless  liquid,  which  can  be  transported  in  steel  cylinders  or 
shipped  like  gasoline  in  tank  cars. 

Chemical  properties  of  sulfur  dioxide.  Sulfur  dioxide  does  not 
burn  in  air,  but  it  may  be  made  to  combine  with  another  atom  of 
oxygen,  forming  sulfur  trioxide,  SO3,  by  passing  a  mixture  of  sulfur 
dioxide  and  air  over  a  catalyst,  such  as  platinum  or  vanadium  pent- 

2SO2  +  O2  -*  2SO3 

This  ability  of  sulfur  dioxide  to  combine  with  more  oxygen  makes 
sulfur  dioxide  a  reducing  agent.  Sulfur  dioxide  may  also  act  as  an 
oxidizing  agent.  When,  for  example,  it  is  passed  through  a  solution 
of  hydrosulfuric  acid,  H2S,  it  precipitates  sulfur.  The  sulfur  (S++++) 
in  the  SO2  takes  electrons  from  the  sulfur  (S — )  of  H2S  and  changes 
to  free  sulfur,  S°. 

S++++O2—  +  2H2+S—  -»  2H2+O—  +  3S°  j 

Hydrosulfuric  acid  acts  here  as  the  reducing  agent.  These  two  actions 
of  sulfur  dioxide  illustrate  the  fact  that  a  substance  may  be  either  an 
oxidizing  agent  or  a  reducing  agent,  depending  upon  the  substance 
to  be  reduced  or  oxidized. 







check  valve- 




•H^lpfc^ expansion  valve 

freezing  tray 

Fig.  69.  An  electric  refrigerator. 
The  refrigerant  may  be  sulfur 
dioxide,  methyl  chloride,  Freon, 


electric  motor  compressor 

When  sulfur  dioxide  dissolves  in  water,  it  combines  with  some  of 
the  water,  forming  sulfurous  acid.  Sulfur  dioxide,  then,  is  an  acid 

SO2  +  H2O  — >  H2SO3  (sulfurous  acid) 

Sulfur  dioxide  has  many  uses.  The  most  common  use  of  sulfur 
dioxide  is  in  the  manufacture  of  sulfuric  acid.  It  is  used  also  to  some 
extent  instead  of  ammonia  in  mechanical  refrigeration.  Although 
liquid  sulfur  dioxide  during  evaporation  absorbs  only  one-third  as 
much  heat  as  ammonia  does,  it  liquefies  under  a  lower  pressure  and 
hence  is  used  in  some  household  refrigerators. 

Homer  in  his  Odyssey  cited  the  use  of  burning  sulfur  in  fumiga- 
tion. Today  rooms,  houses,  or  warehouses  are  sometimes  fumigated 
by  burning  sulfur  candles  or  by  liberating  sulfur  dioxide  gas  from 
cylinders.  However,  as  a  fumigant  sulfur  dioxide  has  been  almost 
wholly  supplanted  by  more  effective  substances.  It  is  even  extremely 
doubtful  that  fumigation  by  sulfur  candles  was  ever  very  effective. 

How  sulfurous  acid,  H,SO3,  is  prepared  in  the  laboratory.  In 
the  laboratory  preparation  of  sulfur  dioxide,  SO2  is  not  the  product 
first  formed.  Sulfurous  acid  is  the  first  product  of  the  reaction,  but 
it>J£  impossible  to  stop  the  reaction  at  this  point. 

Na2SO3  +  H2SO4  -»  Na2SO4  +  H2SO3 

Sulfurous  acid  is  prepared  by  passing  sulfur  dioxide  gas  through 
cold  water. 

H2O  +  SO2  -» H2SO3 


Properties  of  sulfurous  acid.  Sulfurous  acid  is  a  colorless  solution 
with  a  suffocating  odor  (sulfur  dioxide) .  It  is  unstable  and  decom- 
poses readily  into  water  and  sulfur  dioxide.  By  boiling  sulfurous 
acicl,  all  the  sulfur  dioxide  may  be  driven  off.  The  reaction  is  re- 
versible, depending  upon  the  temperature. 

H20  +  SO.  t±  H2SO3 

Sulfurous  acid  is  a  very  weak  acid  because  it  dissociates  only 
slightly,  forming  few  hydrogen  ions.  When  neutralized  with  a  base, 
it  forms  salts  known  as  sulfites. 

H2SO,  +  2NaOH  ->  Na2SO3  (sodium  sulfite)  +  2H2O 

On  exposure  to  air,  sulfurous  acid  may  be  oxidized  to  sulfuric  acid 
to  some  extent. 

2H2SO3  +  O2  -»  2H2SO4 

In  a  similar  way,  sulfites  may  be  oxidized  to  sulfates. 

How  sulfurous  acid,  sulfites,  and  bisulfites  are  used.  The  tend- 
ency of  one  molecule  of  sulfurous  acid  to  unite  with  one  atom  of 
oxygen,  forming  sulfuric  acid,  makes  it  a  valuable  reducing  agent. 
This  property  of  sulfurous  acid  enables  it  to  be  used  in  the  bleach- 
ing of  wood  pulp,  straw,  silk,  feathers,  dried  fruits,  flour,  molasses, 
and  canned  corn  —  all  of  which  would  be  partially  destroyed  by 
other  bleaching  agents. 

Fruits  that  are  to  be  dyed  are  frequently  bleached  with  sulfur 
dioxide  (as  sulfurous  acicl)  .  For  example,  in  making  maraschino 
cherries,  the  cherries  are  first  bleached  and  then  dyed  any  desired 
color  with  edible  dye.  If  sulfurous  acid  is  used  either  in  bleaching 
or  in  preserving  a  food  that  is  sold  in  interstate  commerce,  this  fact 
must  be  clearly  stated  on  the  label  of  the  package  in  which  it  is  sold. 

Chemical  Construction   Corporation 

A  sulfur  combustion  furnace 
for  the  commercial  production 
of  sulfur  dioxide.  Melted  sulfur 
is  pumped  into  the  furnace 
and  burned  to  form  the  gas. 


Substances  of  vegetable  origin  are  not  easily  destroyed  by  sulfurous 
acid,  and  therefore  their  coloring  matter  may  be  reduced  to  colorless 
compounds  by  this  weak  acid.  However,  the  bleaching  action  of 
sulfurous  acid  is  not  as  lasting  as  that  of  chlorine  or  hydrogen  per- 
oxide, because  the  coloring  matter  that  has  been  reduced  may,  on 
exposure  to  air,  become  reoxidized.  This  explains  in  part  why  straw 
hats,  bleached  with  sulfurous  acid,  may  turn  brown. 

Calcium  bisulfite,  Ca  (HSO3)  2,  containing  dissolved  SO2,  is  used  in 
converting  ground  wood  pulp  into  sulftte  pulp  by  dissolving  out  the 
gluelike  substances  found  mixed  with  the  cellulose  of  the  wood  pulp 
(see  page  598) .  Sulfite  pulp  is  used  in  making  paper. 

Sodium  sulftte  mixed  with  a  solid  acid,  such  as  oxalic  (white  crys- 
tals) ,  makes  up  the  dry  powder  used  in  bleaching  straw  hats.  The 
action  of  this  mixture  of  acid  and  sulfite  illustrates  the  part  that  ions 
play  in  chemical  reactions.  While  the  dry  solid  acid  and  sulfite  are 
in  contact,  no  action  occurs.  As  soon  as  water  is  added,  however,  the 
substances  go  into  solution,  dissociate,  and  the  ions  formed  react, 
producing  sulfurous  acid,  which  bleaches.  Na2SO3  is  also  used  as  an 
antichlor  (a  substance  used  to  remove  any  excess  of  chlorine  after 
bleaching)  and  as  an  oxygen-removal  agent  in  treating  boiler-feed 

How  to  test  for  a  sulfite.  The  characteristic  od6r  of  the  sulfur 
dioxide  liberated  when  an  acid  is  added  to  a  sulfite  constitutes  a 
test  for  any  sulfite. 

CaSO3  +  2HC1  ->  CaCl2  +  H2O  +  SO2  T 

Preparation  and  properties  of  sulfur  trioxide.  When  sulfur  di- 
oxide is  mixed  with  air  and  the  mixture  is  passed  over  a  catalyst 
such  as  heated  platinum,  sulfur  trioxide  is  formed. 

2SO2  +  O2  ->  2SO8 

This  constitutes  the  key  reaction  in  the  most  important  commercial 
method  of  manufacturing  sulfuric  acid,  the  contact  process,  so  named 
because  contact  of  the  gases  with  this  platinum  catalyst  plays  so  im- 
portant a  part. 

Sulfur  trioxide  is  a  white  solid,  but  it  melts  at  about  room  temper- 
ature and  boils  at  46 °C.  The  liquid  fumes  in  moist  air  and  reacts 
vigorously  with  water  making  a  hissing  sound,  liberating  much  heat 
and  forming  sulfuric  acid.  Thus,  sulfur  trioxide  is  the  anhydride  of 
sulfuric  acid. 

SO3  +  H2O  -» H2SO4 


,""  for  '"«' 


Monsanto  Chemical  Company  and  Davison  Chemical  Corporation 

A  contact  process  installation.  Trace  the  process  in  this  photograph  and  that  on 
page  311. 

How  sulfuric  acid  is  made  by  the  contact  process.  The  manu- 
facture of  sulfuric  acid  by  the  contact  process  may  be  conveniently 
divided  into  three  stages:  (1)  burning  sulfur  to  sulfur  dioxide  in 
a  combustion  furnace,  (2)  changing  this  sulfur  dioxide  to  sulfur  tri- 
oxide  by  passing  sulfur  dioxide  and  air  over  a  catalyst  in  a  converter, 
(3)  changing  the  sulfur  trioxide  to  sulfuric  acid  by  passing  the  sul- 
fur trioxide  into  absorption  towers  through  which  sulfuric  acid  is 
circulated.  The  chemical  reactions  that  take  place  are: 

1)  S  +  O2  ->  SO2 

2)  2S02  4-  02  ->  2S03 

3)  SO3  4-  H2O  [+ H?SO4!  -»  i2|H2SO4 

1)  The  sulfur  combustion  furnace.  In  the  combustion  furnace, 
sulfur  is  brought  into  intimate  contact  with  thoroughly  dried  air. 
A  blower  supplies  this  air  at  sufficient  pressure  to  force  it  through 
the  furnace  and  to  force  the  sulfur  dioxide  from  the  furnace  through 
the  rest  of  the  system.  The  heat  produced  by  the  combustion  of  the 
sulfur  is  recovered  by  means  of  a  waste-heat  boiler.  Part  of  this  heat 
is  used  in  some  plants  to  melt  the  sulfur  that  is  sprayed  into  the 



2)  The  converter.  The  converter  is  simply  a  large  chamber  con- 
taining many  perforated  shelves  covered  with  a  catalyst  in  whose 
presence  the  union  of  warm  sulfur  dioxide  and  oxygen  takes  place, 
forming  sulfur  trioxide.  Noteworthy  developments  in  the  contact 
process  have  been  the  introduction  of  vanadium  pentoxide  and 
platinized  silica  gel  catalysts  as  substitutes  for  the  more  expensive 
platinum  mass.  The  vanadium  catalyst  has  the  further  advantage  of 
being  insensitive  to  poisoning  by  impurities,  such  as  arsenic,  anti- 
mony, selenium,  and  chlorine,  that  may  be  present  in  the  sulfur. 

Since,  technically,  the  greatest  cost  in  processes  that  depend  upon 
a  catalyst  is  keeping  the  catalyst  entirely  free  from  minute  traces  of 
poisons,  a  catalyst  less  affected  by  such  impurities  is  preferred.  The 
introduction  and  use  of  the  vanadium  catalyst  was  soon  followed  by 
a  platinum  catalyst  that  is  immune  to  arsenic  poisoning  and  may 
be  operated  at  a  lower  temperature. 

The  change  of  sulfur  dioxide  to  sulfur  trioxide,  which  occurs  on 
the  surface  of  the  catalyst,  is  an  exothermic  one,  that  is,  heat  is  given 
off.  Careful  temperature  control  is  necessary  to  insure  highest  yields. 
Below  450° C.,  some  SO2  escapes  oxidation  to  SO3.  Above  450 °C., 
some  SO3  is  decomposed  to  SO2.  At  a  temperature  of  450°C.  the 
reversible  reaction 

2SO2  +  O2  *=±  2SO3 

goes  to  the  right  almost  completely.  Only  about  three  percent  of  the 
total  amount  of  SO3  theoretically  possible  is  changed  back  to  SO2. 

3)  The  absorption  towers.  From  the  converter,  the  sultur  trioxide 
passes  through  a  cooler,  and  then  into  a  series  of  three  or  four  ab- 
sorption towers.  Through  the  first  two  of  these  towers,  filled  with 
quartz  pebbles  or  acid-resisting  packing  rings,  concentrated  sulfuric 
acid  is  circulated.  This  concentrated  sulfuric  acid  absorbs  the  sulfur 
trioxide,  forming  sulfuric  acid  of  more  than  98-percent  concentra- 
tion. SO3  vapor  is  absorbed  more  easily  by  concentrated  H2SO4  than 
by  water.  The  third  tower  is  a  coke-filled  filter  tower,  which  absorbs 
any  acid  vapor  that  might  otherwise  escape  into  the  air. 

A  contact  plant  is  practically  automatic.  Control  of  the  process  is 
maintained  from  one  central  room  containing  all  the  recording  and 
controlling  instruments. 

How  the  contact  process  compares  with  the  lead-chamber  process 
of  preparing  sulfuric  acid.  The  older  commercial  process  for  making 
sulfuric  acid  is  the  lead-chamber  process.  It  was  introduced  in  Bir- 
mingham, England,  in  1746  by  Dr.  Roebuck,  a  physician.  Its  use 



marked  the  beginning  of  chemical  manufacture  on  a  large  scale, 
for  the  Industrial  Revolution  was  just  then  getting  under  way. 

In  recent  years,  the  lead-chamber  process  has  given  way  gradually 
to  the  contact  process.  Because  of  the  large  volumes  of  gases  that 
must  react  (without  a  metallic  catalyst) ,  large  reaction  chambers 
are  needed.  On  account  of  the  action  of  dilute  su  If  uric  acid  on  iron, 
it  is  necessary  to  line  these  large  rooms  with  lead  sheets,  two  inches 
thick  —  a  second  item  of  expense.  Some  of  the  chemical  reactions 
that  occur  in  the  chamber  process  are  still  not  thoroughly  under- 

Sulfur  dioxide,  steam,  and  the  oxygen  of  the  air  are  converted  into 
dilute  sulfuric  acid  by  means  of  the  oxides  of  nitrogen,  which  act  as 
catalytic  agents.  The  change  of  NO  into  NO2,  which  takes  place  in 
air,  and  the  subsequent  reversal  of  this  reaction  are  indirect  means 
of  getting  an  extra  atom  of  oxygen  to  combine  with  sulfur  dioxide. 
At  the  end  of  the  process  a  corrosive,  sour  drizzle  of  dilute  sulfuric 
acid  falls  on  the  floor  of  the  lead  chambers. 

Another  disadvantage  of  the  lead-chamber  process  is  that  only 
dilute  sulfuric  acid,  often  impure  and  of  not  more  than  78-percent 
concentration,  can  be  made.  If  concentrated  acid  is  required,  this 
dilute  acid  must  be  concentrated  by  evaporation  with  heat. 

The  chamber  process  can  compete  with  the  contact  process  only 
if  acid  of  not  greater  than  78-percent  concentration  is  required,  as 
in  the  manufacture  of  phosphate  fertilizers  and  in  the  pickling  of 
steel.  Sulfuric  acid  of  less  than  78-percent  concentration  must  be 
shipped  in  glass  or  in  lead-lined  steel  containers,  while  acid  of 

In  these  towers  the  final  steps  of  the  contact  process  take  place. 

Monsanto   Chemical   Company  and   Garfield   Chemical  and  Manufacturing   Corporation 



Fig.  70.  Diluting  concentrated 
sulfuric  acid.  Why  is  such  care 
not  necessary  when  diluting  hy- 
drochloric acid? 





greater  than  78-percent  concentration  may  be  shipped  in  steel  drums 
or  tank  cars.  Why? 

Physical  properties  of  sulfuric  acid.  Sulfuric  acid,  or  oil  of  vitriol, 
is  a  water  solution  of  hydrogen  sulfate  (a  liquid)  .  Concentrated  sul- 
furic acid  contains  98  percent  acid,  has  a  specific  gravity  of  1.84,  and 
boils  at  338°C.  It  mixes  with  water  in  all  proportions,  liberating 
much  heat.  For  this  reason,  great  care  should  be  used  in  diluting 
concentrated  H,SO4.  The  heavier  sulfuric  acid  should  be  added  to 
cold  water  slowly,  and  the  mixture  constantly  stirred  as  shown  in 
Fig.  70.  Always  add  the  acid  to  the  water.  Oleum,  or  fuming  sulfuric 
acid,  contains  dissolved  sulfur  trioxide,  and  usually  its  formula  is 
written  H,SO4  •  SO3. 

Chemical  properties  of  sulfuric  acid.  A  water  solution  of  sulfuric 
acid  is  a  very  strong  acid,  because  it  dissociates  to  a  high  degree.  The 
concentrated  acid,  because  it  does  not  dissociate  to  any  appreciable 
extent,  is  weaker  than  the  dilute  acid. 

Because  of  its  strong  attraction  for  water,  sulfuric  acid  is  an  excel- 
lent drying,  or  dehydrating,  agent.  Its  attraction  for  water  is  so 
intense  that  sulfuric  acid  will  remove  the  hydrogen  and  oxygen  pres- 
ent in  carbohydrates  in  which  these  elements  occur  in  the  same  pro- 
portion as  in  water.  Thus,  when  concentrated  sulfuric  acid  is  poured 

over  sugar,  C12H,2On,  it  removes  the  H,2OU  as  1  1  molecules  of  steam, 
and  a  black  mass  of  pure  carbon  remains. 

-»  11H2O  +  12C 

Likewise,  wood,  which  is  composed  of  cellulose,  C(1H10O5,  chars 
when  clipped  in  concentrated  sulfuric  acid.  Cotton,  wool,  and  other 
substances  react  similarly.  The  dehydrating  action  of  sulfuric  acid 
accounts  also  for  the  severe  burns  it  may  produce  on  the  skin. 

At  450°C.  and  atmospheric  pressure,  sulfuric  acid  decomposes 
completely  into  water  and  sulfur  trioxide.  This  chemical  reaction 
is  reversible. 

H2SO4  ^  SO3  +  H2O 

It  has  been  pointed  out  that  nitric:  acid  is  an  oxidizing  agent  be- 
cause, when  added  to  a  metal,  it  forms  water  instead  of  hydrogen, 
and  liberates  a  gaseous  oxide.  Warm  concentrated  sulfuric  acid, 
which  behaves  in  a  manner  similar  to  nitric  acid,  is  also  an  excellent 
oxidizing  agent. 

Cu  +  2H2S04  ->  CuS04  +  2H2O  +  SO2  1 

This  reaction  between  concentrated  sulfuric  acid  and  copper  illus- 
trates another  common  laboratory  method  for  preparing  sulfur  di- 
oxide. In  contact  with  certain  metals,  including  zinc,  iron,  and 
magnesium,  dilute  sulfuric  acid  liberates  hydrogen  as  shown  in  the 
following  equation, 

Mg  +  H2SO4  -»  MgSO4  +H2  1 

Iron  and  steel  (pickling) 

Rayon  and  film 

Other  metallurgical 



Sulfuric  acid  aids  agriculture.  In  a  recent  year,  the  United  States 
produced  about  50  percent  of  the  sulfuric  acid  consumed  by  the 
entire  world.  The  domestic  consumption  amounted  to  slightly  more 
than  12  million  tons.  While  it  is  not  likely  that  most  farmers  believe 
themselves  in  any  large  measure  dependent  on  sulfuric  acid,  it  is  a 
fact  that  the  fertilizers  used  very  widely  in  agriculture  account  for 
almost  one-fourth  of  the  sulfuric  acid  consumed  each  year. 

Phosphorus  occurs  in  nature  in  a  fairly  plentiful  mineral  called 
tricalcium  phosphate,  or  rock  phosphate.  Unfortunately,  rock  phos- 
phate is  insoluble  in  water  and  therefore  cannot  be  utilized  by 
plants.  However,  when  treated  with  H2SO4,  rock  phosphate  changes 
into  a  soluble  compound  monocaldum  phosphate  (the  active  ingre- 
dient of  superphosphate  —  see  page  487) .  The  phosphorus  is  then 
available  as  plant  food. 

In  cases  where  the  soil  is  acid,  however,  pulverized  raw  rock  phos- 
phate is  applied  directly.  The  acids  of  the  soil  attack  the  phosphate 
slowly  and  make  part  of  it  available  as  plant  food. 

Sulfuric  acid  in  production  of  petroleum  products.  The  petroleum 
industries  consume  about  one-eighth  of  all  the  sulfuric  acid  produced 
in  the  United  States.  Crude  petroleum  contains  a  large  number  of 
carbon  compounds  that  are  dark  in  color  or  that  become  dark  from 
exposure  to  air.  These  compounds  are  removed  from  crude  oil  by 
treatment  with  sulfuric  acid.  In  this  way  most  of  the  waxy  and 
gummy  materials  that  clog  burners  and  carburetors  are  removed. 

Other  uses  of  sulfuric  acid.  Another  important  use  of  sulfuric  acid 
is  in  the  cleaning,  or  pickling,  of  sheet  steel  before  covering  the  steel 
with  a  layer  of  zinc  in  galvanizing,  or  tin  in  tin-plating.  The  iron 
sulfate  that  is  formed  as  a  byproduct  is  crystallized  and  is  used  in 
making  inks.  The  common  name  for  crystallized  iron  sulfate  is 
copperas,  once  known  as  green  vitriol. 

Fe  +  H2SO4  +  7H2O  -»  H2 1  +  FcSO4  •  7H2O  (copperas) 

The  manufacture  of  explosives  is  dependent  upon  large  quanti- 
ties of  concentrated  sulfuric  acid.  It  is  used  as  a  dehydrating  agent 
in  the  process  of  nitrating  the  many  substances  from  which  explo- 
sives are  made.  Sulfuric  acid  is  used  in  the  production  of  textiles  and 
rayon  and  cellulose  film.  Many  acids,  including  hydrochloric  and 
nitric,  are  made  with  the  aid  of  sulfuric  acid.  Others  of  the  hundreds 
of  processes  in  which  sulfuric  acid  is  used  include  the  making  of  sul- 
fates,  paper,  rayon,  leather,  celluloid,  dyes,  dn$$p,  and  paints.  (Many 
of  these  processes  will  be  described  later.)  Sutewj&acid  is  used  also 
as  the  electrolyte  in  lead  storage  batteries.  ^%r 


Normal,  acid,  and  basic  salts.  Zinc  sulfate,  ZnSO4,  is  formed  by  the 
reaction  of  zinc  and  sulfuric  acid.  Here  both  of  the  hydrogen  ions 
of  the  acid  have  been  replaced  by  the  metal,  and  the  resulting  salt 
is  a  normal  salt,  that  is,  a  salt  in  which  all  the  hydrogen  of  the  acid 
has  been  replaced  by  a  metal.  In  addition  to  forming  normal  salts, 
acids  such  as  H2SO4,  H2COS,  and  H2SO3,  which  contain  two  hydro- 
gen atoms,  can  also  form  acid  salts,  and  are  called  dibasic  acids.  Boric 
acid,  H3BO3,  which  contains  three  replaceable  hydrogen  atoms,  is  a 
tribasic  acid. 

In  the  preparation  of  hydrochloric  acid  by  the  action  of  sulfuric 
acid  on  sodium  chloride,  sodium  bisulfate,  NaHSO4  (also  called  so- 
dium acid  sulfate) ,  is  obtained  as  one  of  the  byproducts. 

NaCl  +  H2SO4  ->  NaHSO4  +  HC1 

Here  only  one  of  the  hydrogen  ions  has  been  replaced,  and  the  other 
remains  in  the  salt,  which  is  called  an  acid  salt.  Thus  it  can  be  seen 
that  either  one  or  both  of  the  hydrogen  ions  of  sulfuric  acid  may  be 
replaced  by  a  metal. 

Acids  such  as  HC1  and  HNOa  contain  only  one  hydrogen  atom  and 
are  called  monobasic  acids.  A  salt  containing  one  or  more  hydroxide 
groups  is  called  a  basic  salt.  Such  a  salt  is  basic  lead  carbonate 
Pb8  (OH)  2  (CO,)  a  (see  page  449) . 

Glauber  salt  and  epsom  salt.  One  of  the  most  famous  chemists  of 
the  seventeenth  century  was  a  Bavarian,  Johann  Glauber.  When  he 
was  21  he  was  attacked  by  a  fever  and  was  advised  to  drink  the  water 
of  a  certain  well.  After  recovering  from  his  illness  he  analyzed  wa- 
ter from  this  well,  and  extracted  from  it  crystals  of  a  salt  which  he 
called  sal  mirabile,  the  wonderful  salt.  He  recommended  it  as  a 
"splendid  medicine  for  internal  and  external  use."  This  salt, 
Na2SO4.  10H2O,  known  as  glauber  salt  has  been  used  as  a  laxative 
for  more  than  300  years.  Anhydrous  Na2SO4,  also  known  as  salt  cake, 
is  used  in  the  manufacture  of  paper  and  glass. 

Soon  after  the  introduction  of  glauber  salt,  an  English  physician, 
Nehemiah  Grew,  who  discovered  that  plants  have  sex,  extracted 
a  salt  from  some  springs  in  the  village  of  Epsom,  near  London, 
England.  He  wrote  a  book  on  the  medicinal  value  of  this  salt, 
MgSO4  •  7H2O,  and  thus  epsom  salt,  a  white,  soluble,  crystalline 
compound  became  a  rival  laxative. 

Many  proprietary  medicines,  mineral-water  crystals,  and  mineral 
waters  sold  today  at  ridiculously  high  prices  are  laxatives  containing 
one  or  both  of  these  salts.  Such  salt  mixtures  and  solutions  are  sold 
to  be  used,  often  dangerously,  as  "the  modern  way  to  slenderize." 


Reducing  preparations  of  this  kind  should  not  be  taken  without  the 
advice  of  a  competent  physician. 

How  to  test  for  a  sulfate  ion.  Most  sulfates  are  soluble  in  water. 
However,  the  sulfates  of  barium,  lead,  and  strontium  are  insoluble. 
When  barium  chloride  is  added  to  a  sulfate,  a  white  precipitate, 
barium  sulfate,  is  formed.  This  white  precipitate  is  distinguished 
from  other  barium  salts  by  the  fact  that  barium  sulfate  is  insoluble 
in  hydrochloric  acid. 

Na2SO4  +  BaCl2  -»  BaSO4 J  +  2NaCl 
This  reaction  goes  to  completion.  Why? 


Aaron,  Harold.  Good  Health  and  Bad  Medicine.  Consumers 
Union  of  United  States,  Inc.,  New  York,  1940.  Carefully  pre- 
pared and  well-written  materials  on  self-medication. 

Goldblatt,  L.  A.,  Ed.  Collateral  Readings  in  Inorganic  Chem- 
istry. D.  Appleton-Century  Co.,  New  York,  2nd  series,  1942. 
Article  15  contains  an  excellent  description  of  a  "Contact 
Sulfuric  Acid  Plant." 

Leicester,  H.  M.,  and  Klickstein,  H.  S.  Source  Book  in  Chem- 
istry. McGraw-Hill  Book  Company,  New  York,  1952.  Pages 
11-16  contain  Georg  Agricola's  description  of  the  manufacture 
of  vitriol,  from  De  re  Metallica. 


Group  A 

1.  Copy  and  complete  the  following  statements.  Do  not 
write  in  this  book.  SO2  is  prepared  for  industrial  use  by  burn- 
ing ...  or  roasting  ....  It  occurs  in  nature  in  small  amounts 
near  ....  The  most  important  use  of  SO2  is  in  the  manufacture 
of  ....  Its  most  characteristic  physical  property  is  .... 

2.  (a)  What  two  methods  are  used  in  preparing  SO2  in 
the  laboratory?  (b)  Which  method  involves  reduction?  (c)  Ex- 

3.  (a)  Make  a  labeled  diagram  showing  how  SO2  is  pre- 
pared from  Na2SO3.   (b)  Write  the  equation  for  the  reaction 
that  occurs. 

4.  Compare  the  water  solubility  of  SOa,*faH3,  HC1,  and  O2. 


5.  Write  three  equations  illustrating  three  chemical  prop- 
erties of  SO2. 

6.  Describe  three  uses  of  SO 


7.  For  what  purpose  is  SO2  used  in  the  household? 

8.  (a)  Is  SO2  formed  first  by   the  action  of  H2SO4  on 
Na2SO3?  (b)  Explain. 

9.  (a)  Write    the    reversible    reaction   between   SO2    and 
H2O.   (b)  How  can  the  reaction  be  made  to  go  to  the  right? 
(c)  to  the  left? 

10.    (a)  Why  is  H2SO3  a  weak  acid?  (b)  What  is  a  sulfite? 
1  1  .  Write  an  equation  showing  the  formation  of  Na2SO3. 

12.  (a)  By    what    chemical    process    does    H2SO3    bleach? 
(b)  Into  what  does  H2SOS  change  as  it  bleaches?  (c)  Compare 
the  way  in  which  H2SO3  bleaches  with  the  way  in  which  C12 

13.  Write  H2SO3,  H2O2,  and  C12  in  a  column.  Opposite  each 
write  the  names  of  the  substances  that  it  is  best  suited  to 

14-    (a)  Why  is  bleaching  with  H2SO3  less  permanent  than 
bleaching  with  other  chemicals?  (b)  Illustrate. 

15.  Distinguish  between  a  sulfite,  a  sulfate,  and  a  sulfide. 

16.  How  much  S  would  be  needed  to  prepare  500  tons  of 
98  percent  H2SO4? 

17.  Cu  is  added  to  concentrated  H2SO4  and  448  ml.  of  SO2 
are  liberated.  How  much  H2SO4  is  decomposed? 

18.  Calculate  the  percentage  of  S  in  each  of  its  two  oxides. 

19.  What  are  three  properties  of  SO3? 

20.  Write  the  equation  for  the  burning  (roasting)  of  FeS2. 

21.  What  are  the  stages  in  the  contact  process  of  making 

22.  Write  equations   that  show  the   three  chief  chemical 
changes  that  take  place  in  the  preparation  of  H2SO4  by  the 
contact  process. 

23.  (a)  Why  must  the  catalyst  used  in  the  contact  process 
be  chosen  with  special  care?    (b)  What  advantage  has  V2O5 
over  the  old  plantinum  catalyst  used  in  the  manufacture  of 


24.  Describe  briefly  the  chemical  change  that  is  the  basis 
of  the  chamber  process  for  making  H2SO4. 

25.  (a)  What    are    the    disadvantages    of    manufacturing 
H2SO4  by  the  chamber  process?    (b)  This  process  competes 
favorably  with  the  contact  process  in  producing  H2SO4  for 
what  industry?  (c)  Why? 

".  i " " " ; 

26.  What  are  the  physical  properties  of  oil  of  vitriol? 

27.  In  mixing  concentrated  H2SO4  and  H2O,  why  should 
the  acid  always  be  added  to  the  H2O  rather  than  the  H2O  to 
the  acid? 

28.  Devise  an  experiment  to  show  the  dehydrating  action  of 

29.  By  its  action  on  Cu,  show  that  concentrated  H2SO4  is 
an  oxidizing  agent. 

30.  (a)  Make  a  list  of  the  uses  of  H2SO4.    (b)  From  the 
viewpoint  of  the  amount  consumed,  what  is  its  chief  use? 
(c)  Explain  fully,  using  an  equation  to  show  the  reaction  in- 

31.  Why  does  the  refining  of  petroleum  require  immense 
quantities  of  H2SO4? 

32.  Name  three  experiments  in  which  you  used  H2SO4  in 
making  laboratory  preparations. 

33.  What  is  sulfite  pulp? 

34.  What  is  the  difference  between  a  normal,  a  basic,  and  an 
acid  salt? 

35.  (a)  Name  three  sulfates  and  (b)  state  an  important  use 
of  each. 

Group  B 

36.  Why  is  SO3  not  added  directly  to  water  in  making 
H2SO4  in  the  last  stage  of  the  contact  process? 

37.  Make  a  diagram  of  the  apparatus  you  would  use  in  pre- 
paring SO2  in  the  laboratory  by  the  action  of  H2SO4  on  Cu. 

38.  Explain:   (1)  Black  rings  often  form  on  wooden  shelves 
that  hold  bottles  of  H2SO4.  (2)  A  full  battle  of  concentrated 
H2SO4  overflows  when  exposed  to  air.   (3)  An  open  bottle  of 
concentrated   H2SO4   is   sometimes   placed   in   a   desiccator. 
(4)  Frightful  burns  result  from  getting  concentrated  H2SO4  on 
the  hands.  * 


39.  Compare  the  three  strong  acids  with  respect  to  their 
action  on  Cu  and  Zn.  Consider  both  dilute  and  concentrated 

40.  A  load  of  scrap  iron  weighing  2715  pounds  and  con- 
taining 95.5%  iron  is  added  to  a  large  vat  of  dilute  sulfuric 
acid.  What  is  the  maximum  amount  of  ferrous  sulfate  that  can 
be  obtained? 


1.  If  you  happen  to  have  epsom  salt  in  your  medicine  chest, 
bring  it  to  school,  read  the  label  to  your  class  and  make  a  test 
for  a  sulfate  with  some  barium  chloride  solution  your  teacher 
will  supply. 

2.  Write  a  two-  or  three-page  report  to  justify  the  statement 
that  H2SO4  is  "king  of  chemicals,"   and  the  "basis  of  the 
machine  age."  Prepare  a  debate  or  a  class  discussion  based  on 
these  statements. 




.  .  .  It  (the  light  from  my  lamp) 
was  hurled  out  of  the  sun  200,000,- 
ooo  years  ago,  and  was  captured  by 
the  leaves  of  the  Carboniferous  tree- 
fern  forests,  fell  with  the  falling 
plant,  was  buried,  fossilized,  dug  up 
and  resurrected.  It  is  the  same  light. 
And,  in  my  little  fig  tree  as  in  the 
ancient  ferns,  it  is  the  same  unchang- 
ing green  stuff  from  age  to  age. 
Donald  Culross  Peattie  in  Flowering 
Earth,  1939. 

The  fascinating  chemistry  of  carbon.  Carbon  is  the  eleventh  most 
abundant  element  of  the  earth  by  weight.  In  the  form  ot  coal  and 
graphite,  it  helps  to  drive  as  well  as  lubricate  the  wheels  of  our 
machine  age.  As  petroleum,  it  propels  our  locomotives,  ships,  auto- 
mobiles, and  airplanes.  In  the  form  of  diamond  and  "Carborundum" 
(synthesized  directly  from  free  carbon) ,  carbon  is  used  in  making 
tools  that  drill  our  hardest  rocks  and  grind  and  polish  our  machines, 
tools,  and  household  utensils. 

In  combined  carbon  we  find  one  of  the  most  fascinating  stories 
in  the  whole  romance  of  chemistry.  Each  year  carbon's  650,000 
known  compounds  are  being  increased  in  number  by  thousands  of 
compounds  synthesized  in  our  industrial  and  university  laborato- 
ries. Synthetic  chemistry  reaches  its  greatest  development  in  organic 
chemistry,  the  chemistry  of  the  compounds  ot  carbon. 

Woehler  (vu-ler)  ,  in  1828,  synthesized  urea,  the  first  organic  com- 
pound made  outside  the  living  body.  By  this  synthesis  the  old  idea 
that  organic  compounds  could  be  formed  only  in  living  matter  was 
destroyed,  and  a  new  era  in  chemistry  was  ushered  in,  witnessing 
"the  great  tragedy  of  science,  the  slaying  of  a  beautiful  hypothesis 
by  an  ugly  fact." 




One  hundred  and  twenty-five  years  after  Woehler's  achievement, 
sugar  (C^H^O^)  was  synthesized  from  its  elements.  What  a  century 
and  a  quarter  of  research  between  Woehler's  urea  and  synthetic 
cane  sugar!  Berthollet  went  to  the  red  ant  and  learned  the  secret  of 
preparing  formic  acid,  the  liquid  responsible  for  the  sting  of  this  in- 
sect. William  Perkin,  washing  bottles  in  a  London  laboratory,  mixed 
at  random  the  contents  of  two  bottles  and  discovered  a  method  of 
synthesizing  mauve  —  the  first  of  a  long  series  of  organic  dyes  that 
rival  the  colors  of  nature.  And  this  is  only  the  beginning.  The  mind 
fairly  reels  at  the  accomplishments  and  possibilities  of  this  branch  of 
chemistry.  More  of  this  entrancing  story  is  told  in  later  chapters. 

What  is  the  valence  of  carbon?  Carbon  has  an  atomic  weight  of 
12.  The  nucleus  of  the  carbon  atom  contains  six  protons;  hence,  the 
outer  ring  contains  only  four  electrons.  Therefore,  the  carbon  atom 
must  either  lose  four  electrons  or  gain  four  electrons  to  complete  its 
outer  ring.  The  valence  of  carbon  then  is  either  plus  four  or  minus 
four,  depending  upon  the  substance  with  which  it  reacts.  That  is, 
it  may  act  as  either  a  metal  or  a  nonmetal.  Carbon  almost  always 
forms  firm  covalent  bonds  involving  four  pairs  of  shared  electrons. 

The  properties  of  carbon.  Carbon  occurs  in  three  allotropic  forms: 
two  of  them,  diamond  and  graphite,  are  crystalline;  the  third  is 
amorphous,  or  noncrystalline,  in  form.  All  forms  of  carbon  combine 
with  oxygen,  forming  either  carbon  monoxide,  CO,  or  carbon  di- 
oxide, CO2.  The  different  forms  of  carbon  vary  widely  in  the  ease 
with  which  they  may  be  oxidized.  Coal  burns  easily  particularly 
when  powdered.  On  the  other  hand,  diamond  combines  with  oxygen 
only  at  extremely  high  temperatures. 

C  +  O2  — >  CO2  (carbon  dioxide) 
2C  +  O2  — >  2CO  (carbon  monoxide) 

Its  attraction  for  oxygen  makes  carbon  a  good  reducing  agent.  The 
reducing  action  of  carbon  may  be  shown  by  heating  a  mixture  of 


Fig.   71.    Blowpipe   analysis 
of  an  ore  (left)  using  the  re- 
ducing   power   of    charcoal, 
(right)  The  carbon  atom. 


carbon  and  lead  oxide  on  a  charcoal  block,  using  a  blowpipe.  A  lit- 
tle ball  of  pure  lustrous  lead  is  formed  and  CO2  is  given  off. 

2C  4-  Pb3O4  ->  3Pb  +  2CO2  T 

Millions  of  tons  of  coke   (carbon)   are  used  industrially  every  year 
in  reducing  iron  oxide  to  commercial  forms  of  iron. 

Fc2O3  +  3C  ->  2Fc  +  SCO  | 

This  reaction  shows  carbon  acting  as  a  metal,  a  lender  of  electrons. 
Carbon  may  also  behave  as  a  nonmetal,  a  borrower  of  electrons.  At 
the  very  high  temperatures  reached  in  an  electric  furnace,  carbon 
unites  with  metals  such  as  calcium  and  wolfram  and  with  less  electro- 
negative elements  such  as  silicon,  forming  carbides. 

CaO        +3C->          CaC2         +  CO  t 

calcium  oxide  calcium  carbide 

SiO2        +3C-»        SiC        +2COT 

silicon  dioxide  silicon  carbide 

Carbon  unites  directly  with  hydrogen  also  at  the  high  temperatures 
of  an  electric  arc,  forming  acetylene  gas,  C..H.,. 

2C  +  H2  ->  C2H2  T 

All  forms  of  carbon  are  insoluble  in  water,  acids,  and  bases. 
Amorphous  carbon,  however,  is  soluble  in  some  molten  metals  in- 
cluding iron. 

Physically,  the  allotropic  forms  of  carbon  resemble  one  another 
in  some  respects.  They  are  all  odorless,  tasteless,  and  insoluble  in 
water.  On  heating,  they  do  not  melt  but  vaporize. 

The  hardest  substance  known.  In  the  eighteenth  century,  diamond 
was  shown  by  Lavoisier  to  be  almost  pure  carbon.  This  French  chem- 
ist burned  a  diamond  in  the  presence  of  a  distinguished  audience  by 
concentrating  the  sun's  heat  on  it  with  a  lens.  He  showed  that  the 
gas  which  was  formed  was  pure  carbon  dioxide.  This  startling  dis- 
covery strengthened  Lavoisier's  theory  of  burning.  Why? 

Diamond,  chief  among  precious  stones,  is  found  mainly  in  South 
Africa,  Brazil,  and  Borneo.  It  occurs  in  various  sizes  and  states  of 
purity,  buried  deep  in  the  ground  frequently  in  the  pipes  of  extinct 
volcanoes,  or  in  the  loose  sands  along  certain  streams.  The  brilliance 
of  diamond,  caused  by  its  power  of  refracting  light,  is  brought  out 
by  removing  any  surface  impurities  or  imperfectiqns  and  by  cutting 
facets  on  its  surface.  Small  amounts  of  metallic  oxides  in  diamonds 



Shell  Oil  Company  Baunigold  Bros.,  Inc. 

This  asphalt  core,  neatly  cut  from  a  Before  use  in  jewelry,  a  diamond  is 
roadbed  by  a  diamond  drill,  will  be  polished  against  a  rotating  iron  disc 
taken  to  a  laboratory  for  study.  coated  with  olive  oil  and  diamond  dust. 

give  them  various  colors,  the  rarest  of  which  are  red,  green,  and  blue. 
The  density  of  diamond  is  3.5. 

Diamond  is  weighed  in  carats,  a  unit  of  weight  based  on  the  weight 
of  the  bean  of  the  carob  tree.  One  carat  is  equal  to  one-fifth  of  a 
gram.  The  largest  diamond  ever  mined  was  the  flawless  Cullinan 
found  in  South  Africa  in  1905.  It  weighed  3026  carats,  or  21.35 
ounces,  and  was  about  the  size  of  a  man's  fist. 

From  the  first  little  diamond  found  in  a  pebble  picked  up  by  a 
child  in  1867,  the  diamond  industry  of  South  Africa  has  grown  to 
world  leadership.  In  a  recent  year,  more  than  three  million  carats  of 
diamonds  were  mined.  About  half  of  this  amount  is  used  in  jewelry. 

Though  brittle,  diamond  is  the  hardest  substance  known,  and  for 
this  reason  is  used  extensively  in  making  tools  for  drilling  rock, 
grinding,  sawing,  engraving,  and  polishing/ Impure  black  diamonds, 
called  carbonado  and  bort,  are  used  in  drills.)  Diamonds  are  also  used 
as  bearings  in  watches  and  delicate  instruments,  and  in  wire-drawing 
dies  made  by  drilling  tapered  holes  through  diamond  crystals. 

Can  diamonds  be  made  synthetically?  Scientists  believe  that 
diamonds  were  formed  under  great  heat  and  pressure  by  the  crystal- 
lization of  carbon  dissolved  in  molten  iron  or  other  substances.  To 
test  this  theory  of  the  formation  of  diamonds,  Henri  Moissan  (mwa'- 
sa'N') ,  in  1893,  subjected  some  pure  carbon  to  extremely  high  tem- 
peratures and  pressures.  He  dissolved  carbon  in  molten  iron,  and 
quickly  plunged  the  molten  mass  into  cold  water.  The  sudden 
contraction  of  the  iron  exerted  tremendous  pressure  on  the  dissolved 


carbon,  causing  it  to  crystallize.  The  iron  was  dissolved  with  acids, 
and,  aided  by  a  microscope,  Moissan  reported  seeing  synthetic  dia- 
monds formed  of  tiny  cubical  crystals.  Other  attempts  have  been 
made  to  synthesize  diamonds,  but  thus  far  these  attempts  have  not 
proved  a  threat  to  the  diamond  industry. 

Black,  soft,  and  greasy.  A  second  allotropic  form  of  carbon, 
graphite,  is  lustrous,  black,  soft,  and  oily  to  the  touch.  The  great 
difference  in  hardness  between  the  allotropic  forms  of  carbon  is 
caused  by  the  different  arrangement  of  the  atoms.  The  atoms  of  the 
diamond  molecule  are  very  closely  interlaced  and  bound  firmly  to- 
gether. Graphite,  on  the  other  hand,  is  made  up  of  piles  of  flat  lay- 
ers of  atoms  firmly  bound  to  one  another,  but  far  apart  and  only 
feebly  hanging  on  to  the  next  layer.  For  this  reason  graphite  is  soft 
and  slippery.  The  flat  planes  slide  over  one  another  like  cards  in  a 
pack.  A  few  electrons  between  the  planes  are  loosely  held,  making 
graphite  an  excellent  conductor,  whereas  diamond  is  a  nonconductor. 

Natural  graphite,  which  is  found  chiefly  in  Ceylon,  Siberia, 
and  Canada,  contains  many  impurities.  It  is  believed  to  have  been 
formed,  like  diamond,  under  high  temperature  and  pressure. 

Electric  furnaces  and  uses  of  synthetic  graphite.  In  1896  Edward 
Acheson,  ot  Pennsylvania,  patented  a  process  for  making  synthetic 
graphite,,  today  a  widely  used  article  of  commerce.  It  is  superior  to 
natural  graphite,  because  it  is  free  from  grit  and  contains  almost  no 
impurities.  Acheson  was  at  different  times  a  blast-furnace  helper,  oil 
gager,  raih*oad  ticket  agent,  store  clerk,  miner,  bookkeeper,  and  en- 
gineer. He  also  worked  for  Edison  and  was  sent  abroad  to  the  Inter- 
national Exposition  in  Paris  to  exhibit  the  Edison  electric  lamp 
which  he  had  helped  to  improve  with  a  graphite  filament.  His 
method  of  making  synthetic  graphite  consists  of  subjecting  anthra- 
cite coal  or  molded  carbon  to  the  high  temperature  of  an  electric 
furnace,  keeping  air  out  by  covering  the  coal  with  sand. 

An  electric  furnace,  as  you  know,  transforms  electric  energy  into 
heat  energy.  The  resistance  to  an  electric  current  offered  by  loose 
carbon,  in  the  center  ot  the  resistance  type  of  electric  furnace  used 
in  making  graphite,  produces  high  temperatures.  Other  electric  fur- 
naces utilize  the  heat  produced  by  the  passage  of  electricity  through 

carbon  atoms 

Fig.  72.  Diagram  of  a  graphite 
particle  showing  the  layers  of 


Alloy  steels  are  produced 
in  this  electric  furnace.  The 
electrodes  of  the  furnace 
are  made  from  graphite. 

Electro  Metallurgical  L'o  mint  try 

high-resistance  wire,  such  as  an  alloy  of  nickel  and  chromium.  An- 
other type  of  electric  furnace  is  the  arc  type.  Here  the  passage  of 
electricity  across  an  air  gap  produces  a  high  temperature  as  a  result 
of  the  resistance  of  the  air.  Arc-type  electric  furnaces  are  used  in  the 
synthesis  of  nitric  oxide,  NO,  during  the  preparation  of  nitric  acid 
by  nitrogen-fixation. 

Like  other  forms  of  carbon,  graphite  does  not  melt  and  has  an  ex- 
tremely high  point  of  vaporization.  For  this  reason  it  is  used  in  mak- 
ing graphite  crucibles  in  which  steel  and  other  alloys  of  high  melting 
point  are  prepared.  Because  of  its  smoothness,  oilincss,  and  high  tem- 
perature of  fusion,  graphite  is  mixed  with  oils  used  in  lubricating 
heavy,  swiftly  moving  parts  of  machinery,  in  which  even  heavy  oil 
would  evaporate  or  burn  off. 

Graphite  electrodes  are  used  in  electric  furnaces,  arc  lamps,  and 
in  general,  for  conducting  electricity  at  very  high  temperatures. 
Powdered  graphite  is  used  in  dry  cells,  and  in  certain  oil-retaining 
bearings  produced  by  powder  metallurgy  (see  page  467)  . 

Graphite  is  used  also  in  lead  pencils.  The  "lead"  in  a  pencil  is 
graphite  mixed  with  clay  to  vary  the  hardness  of  the  lead.  Graphite 
is  also  a  constituent  of  stove  polishes.  Most  synthetic  graphite  is  pro- 
duced at  Niagara  Falls,  where  hydroelectric  power  is  available  at 
low  cost. 


C0dlv  :|  y  Fig    73    Cross  section  of  an 

electric  furnace  used  in  the 
commercial  production  of 
synthetic  graphite. 

graphite  electrodes 



"Formless,"  yet  of  many  forms.  The  noncrystalline  allotropic  form 
of  carbon  is  amorphous  carbon.  The  purest  form  of  amorphous  car- 
bon can  be  prepared  easily  by  heating  pure  table  sugar.  The  hydro- 
gen and  oxygen  are  driven  off  as  water;  pure  amorphous  carbon  is 

Ci2H22Oii  ->  11H2O  +  12C 

Among  the  most  common  forms  of  amorphous  carbon  are  (a)  lamp- 
black and  carbon  black,  (b)  gas  carbon,  (c)  coke,  (d)  charcoal,  and 
(e)  animal  charcoal,  or  boneblack. 

Valuable  soot.  Lampblack  was  known  to  the  Chinese  and  Egyp- 
tians who  used  it  in  making  ink.  It  is  amorphous  carbon  dust  formed 
by  the  incomplete  burning  of  compounds  rich  in  carbon.  It  is  man- 
ufactured commercially  by  burning  heavy  liquid  hydrocarbons  (com- 
pounds of  hydrogen  and  carbon)  in  an  insufficient  supply  of  oxy- 
gen.vThe  escaping  particles  of  carbon  dust,  or  soot,  from  the  small 
flames  are  collected  on  a  revolving  metal  plate  that  is  kept  at  a  low 
temperature  by  flowing  water.  The  smokiness  of  a  kerosene  lamp 
burning  with  an  insufficient  air  supply  or  with  the  oil  supplied  too 
rapidly  illustrates  the  way  in  which  lampblack  is  formed. 

Carbon  black  is  made  in  the  same  way  as  lampblack  except  that 
natural  gas  is  used  instead  of  liquid  hydrocarbons.  It* is  finer  than 
lampblack,  which  it  is  gradually  replacing.  Carbon  black  is  such  an 
important  item  of  commerce  that  in  a  recent  year  its  production  ac- 
counted for  more  than  ten  percent  of  all  the  natural  gas  produced  in 
the  United  States.  Most  of  the  carbon  black  manufactured  in  this 
country  is  added  to  rubber  tires  to  improve  their  resistance  to  tear 
and  abrasion.  It  is  also  used  as  a  basic  raw  material  in  making  print- 
er's ink,  typewriter  ribbons,  carbon  paper,  phonograph  records,  and 
black  paints. 

Almost  graphite.  Gas  carbon  is  amorphous  carbon  that  collects  on 
the  walls  of  retorts  during  the  manufacture  of  illuminating  gas  from 
coal.  In  structure,  gas  carbon  is  semicrystalline  and  resembles  graph- 
ite. It  is  used  in  the  rods  of  arc  lamps  and  other  electrodes. 

Coke  from  the  destructive  distillation  of  coaL  Heating  a  complex 
carbon  compound  in  an  oven  from  which  air  is*  excluded,  and  con- 
densing the  vapors  formed  is  called  destructive  distillation.  When 
soft  coal  is  so  treated  the  most  important  solid  product  is  coke,  a 
steel-gray,  hard,  brittle  substance  used  chiefly  as  a  reducing  agent 
in  the  extraction  of  iron,  and  as  a  fuel.  The  coal  is  heated  to  2000  °F. 
for  16  hours,  and  then  the  white-hot  coke  is  pushed  out  from  the 
oven  and  quickly  quenched  in  cold  water. 


In  the  byproduct  coke  oven  many  valuable  chemicals  are  recov- 
ered. As  the  gases  pass  through  oil-absorption  tanks,  a  light  oil  sepa- 
rates, from  which  benzene,  toluene,  xylene,  and  naphthalene  are  dis- 
tilled. From  the  coal  tar  which  is  formed  anthracene,  phenol,  and 
pitch  are  also  separated  by  fractional  distillation.  Ammonia  gas  and 
coal  gas  are  also  obtained  (see  pages  362-363  for  a  detailed  account 
of  this  process) .  Less  than  ten  percent  of  our  coke  is  still  made  in 
beehive  ovens  from  which  none  of  these  byproducts  are  recovered. 

Charcoal-broiled.  If  wood  is  subjected  to  destructive  distillation, 
charcoal  is  one  of  the  products.  Charcoal  retains  the  fibrous  struc- 
ture of  wood. 

Powdered  charcoal  is  frequently  put  together  under  tremendous 
pressure,  forming  compact,  pillow-shaped  blocks  called  charcoal 
briquets.  These  burn  with  no  smoke  and  leave  very  little  ash.  They 
are  used  in  camping  or  picnicking,  on  pleasure  boats,  and  in  found- 

Charcoal-broiled  meats,  especially  steaks  and  chops,  are  particu- 
larly delicious.  Meats  broiled  over  charcoal  do  not  pick  up  odors  or 
flavors  from  the  burning  fuel. 

Activated  charcoal.  This  is  a  specially-prepared  charcoal  used  for 
removing  or  adsorbing  gases.  Adsorption  is  caused  by  the  collection 
of  thin  layers  of  molecules  of  gases  and  other  impurities  on  the  sur- 
faces of  the  porous  charcoal.  It  is  not  a  chemical  union  but  a  physical 
attraction.  Ill-smelling  colored  liquids,  passed  through  such  charcoal, 
are  cleared  of  impurities  responsible  for  the  odor  and  color.  Adsorp- 
tion is  used  also  in  recovering  many  industrial  solvents  and  waste  sub- 
stances, in  air  purification,  and  to  concentrate  a  desired  substance 
such  as  the  drug  streptomycin  from  its  mold  culture. 

Activated  charcoal  is  used  in  gas  masks  to  adsorb  poison  gases  met 
in  industry.  "Nuchar"  is  the  trademark  for  an  activated  carbon  of 
vegetable  origin  containing  about  90  percent  pure  carbon. 

A  black  decolorizer.  The  chief  solid  left  after  the  destructive 
distillation  of  clean  cattle  bones  is  a  black  powder  called  boneblack. 
This  boneblack,  or  animal  charcoal,  is  a  mixture  containing  only 
about  ten  percent  amorphous  carbon  and  80  percent  calcium  phos- 
phate, the  chief  compound  found  in  bones.  It,  too,  has  great  adsorp- 
tive  powers  and  is  used  mainly  in  the  refining  of  crude  sugar,  from 
which  it  removes  impurities  that  cause  objectionable  colors  and 
odors.  It  is  used  also  as  a  black  pigment  in  paints. 

How  coal  was  formed.  Coal  is  found  in  all  continents,  including 
Antarctica,  but  the  largest  coal-producing  areas  are  in  North  Amer- 
ica and  Asia.  Scientists  now  believe  that  many  millions  of  years  ago, 


during  the  Carboniferous  period,  great  portions  of  the  earth  were 
covered  with  a  dense  vegetation  more  luxuriant  than  that  found 
today  even  in  tropical  jungles.  As  the  level  of  the  earth  sank  during 
one  of  the  many  upheavals  that  occurred,  vast  portions  of  these 
jungles  were  flooded  by  the  ocean  and  became  swamps.  Later  these 
swamps  were  completely  submerged  and  then  gradually  covered  with 
mud,  sand,  or  clay  deposited  as  sediment  by  streams  and  rivers. 

Partial  decomposition  of  the  wood  and  other  vegetable  matter, 
aided  by  bacteria  and  fungi,  changed  the  woody  material  first  into 
peat.  Minnesota,  Florida,  and  several  other  states  contain  large  peat 
deposits.  Near  Chester,  Wisconsin,  is  a  peat  deposit  that  covers 
32,000  acres.  Its  average  depth  is  six  feet  and  it  contains  about 
40,000  tons  of  air-dried  peat.  Peat  bogs  also  occur  in  Ireland  and  else- 
where. When  dried  and  pressed,  peat  may  be  used  as  a  fuel.  It  burns 
with  a  great  deal  of  smoke  and  produces  little  heat. 

Scientists  believe  that  the  next  step  in  the  formation  of  coal  was 
the  changing  of  the  peat  into  brown  or  black  lignite^  which  still  re- 
tains the  structure  of  the  plants  from  which  it  was  formed.  Some- 
times branches  or  twigs  are  found  in  lignite  in  the  form  in  which 
they  grew.  Lignite  is  found  in  some  of  the  states  west  of  the  Missis- 
sippi, and,  where  no  other  form  of  coal  is  available,  is  used  as  a 
household  fuel.  In  North  Dakota,  for  example,  lignite  is  being 
mined  at  the  present  time  and  utilized  in  making  gas  and  briquet 
fuels,  and  for  the  production  of  electric  power. 

Further  decomposition  and  pressure,  away  from  air,  drove  out 
more  oxygen  and  hydrogen  from  the  lignite,  leaving  hydrocarbons 
and  some  free  carbon  and  forming  bituminous,  or  soft,  coal.  Further 
heat  and  greater  pressures  changed  the  soft  coal  to  jet  black,  lus- 
trous hard  coal,  or  anthracite. 

Strong  evidence  supports  this  theory  of  coal  formation.  The  fos- 
sils embedded  in  coal  deposits  include  forms  of  both  animal  and 
vegetable  life.  Some  of  the  animals  appear  to  have  been  marine. 
Among  the  plants,  mosses  and  ferns  occur  most  commonly. 

The  chief  kinds  of  coal.  The  chief  kinds  of  coal  used  in  this 
country  are  soft,  or  bituminous,  coal,  and  hard  coal,  or  anthracite. 
Most  of  the  anthracite  mined  in  the  United  States  comes  from  Penn- 
sylvania, yet  this  state,  together  with  West  Virginia,  produces  more 
than  half  the  bituminous  coal  mined  in  this  country.  For  every  ton 
of  anthracite  mined,  11  tons  of  bituminous  coal  are  dug.  Anthra- 
cite contains  practically  no  free  carbon  and  produces  less  volatile 
material  than  bituminous  coal  (see  Table  11) .  It  is  cleaner  to  han- 
dle and  burns  with  a  short,  pale  blue  flame. 


The  use  of  anthracite  is  confined  largely  to  the  northeastern  states, 
because  these  states  have  the  advantage  of  being  near  the  deposits 
and  thus  have  short  and  low-cost  transportation  from  the  mines.  In 
other  parts  of  the  country  anthracite  is  relatively  unfamiliar. 

The  problem  of  smog.  Smog,  a  mixture  of  smoke  and  fog,  is 
formed  largely  by  the  incomplete  combustion  of  coal,  fuel  oil,  gas 
and  rubbish.  The  products  of  this  incomplete  combustion  are  oxi- 
dized by  the  o/one  in  the  air  to  create  the  smog.  Smog  may  be  aggra- 
vated by  natural  particles  in  the  air,  such  as  dust  or  pollen,  or  even 
by  local  weather  conditions. 

People  who  live  in  cities  where  smog  is  a  problem  spend  far  more 
for  soaps  and  other  cleansers  than  might  be  necessary  if  better  smoke 
control  were  practiced.  It  has  been  shown  in  several  cities  that  smog 
may  be  eliminated.  However,  success  in  reducing  smoke  depends 
upon  the  cooperation  of  every  consumer  of  fuel  from  the  largest 
industry  to  the  individual  homeowner. 

Smoke  control  is  particularly  necessary  in  cities  in  which  bitumi- 
nous coal  is  a  major  fuel.  Bituminous  coal,  when  burned  in  an  in- 
adequate supply  of  oxygen,  produces  a  very  smoky  flame.  If  the  coal 
is  properly  fired  (sec  page  35!)) ,  smoke  may  be  greatly  reduced. 

British  thermal  units.  A  chief  factor  in  determining  the  price  of 
coal  is  its  fuel  value.  This  value  is  measured  either  in  calories  or  in 
British  thermal  units,  Btu.  One  Btu  is  the  amount  of  heat  required 
to  raise  the  temperature  of  one  pound  of  water  one  degree  Fahren- 
heit. One  Btu  is  equivalent  to  252  calories.  Burning  one  pound  of 
average  anthracite  liberates  about  12,700  Btu;  burning  the  same 
weight  of  bituminous  coal  liberates  about  13,100  Btu.  The  per- 
centage of  ash  and  free  sulfur  is  another  factor  important  in  deter- 
mining price.  In  general,  anthracite  contains  a  smaller  percentage 
of  these  materials  than  does  bituminous  coal. 

A  scene  in  Pittsburgh,  Pa.,  before  (left)  and  after  a  smog-control  program  was 

Allegheny  Conference  on  Community  Development 

TABLE  11.  fl 




Variations  in  fixed  || 

carbon,  volatile  matter,  ^ 

and  moisture  on  an  ash  ^ 

free  basis  of  the  several  || 

ranks  of  coal  produced  ^ 

in  the  United  States  || 



Subbituminous  C 

Subbituminouj  B 

Subbituminous  A 

High -Volatile  C  Bituminous 

High -Volatile  B  Bituminous 

High -Volatile  A  Bituminous 

Medium -Volatile  Bituminous 

Low -Volatile  Bituminous 



Meta- anthracite 




Coal,  a  major  source  of  power.  Coal  has  been  used  as  fuel  for 
centuries.  Marco  Polo,  700  years  ago,  reported  the  Chinese  burning 
"black  stone  dug  out  of  mountains  where  it  runs  in  veins."  But  it  was 
not  widely  used  until  the  sixteenth  century.  By  1661  "a  hellish  and 
dismall  cloud  of  Sea-Coale  hung  perpetually  over  London,"  where 
coal  was  used  to  make  iron,  glass,  and  other  products.  During  the 
Industrial  Revolution  it  became  a  vital  source  of  energy. 

Our  civilization  uses  energy  at  a  staggering  rate.  In  1900,  the  total 
supply  of  energy  from  mineral  fuels  and  water  power  in  the  United 
States  was  7.4  quadrillion  Btu;  by  1950,  the  total  had  reached  36.2 
quadrillion  Btu  —  an  increase  of  almost  400  percent!  Each  year  our 
growing  population  and  industries  need  added  amounts  of  energy. 
It  is  expected  that  before  the  end  of  the  century  the  annual  produc- 
tion of  energy  will  be  more  than  double  the  present  rate. 

Today,  petroleum,  coal,  natural  gas,  and  hydroelectric  dams  are 
our  major  sources  of  energy.  In  a  recent  year,  petroleum  supplied  40 
percent  of  our  nation's  total  energy  requirement;  coal  supplied  34 
percent;  natural  gas,  22  percent;  and  water  power,  4  percent.  In 
the  years  ahead,  the  use  of  nuclear  energy  will,  in  all  probability, 
change  these  figures  radically. 

Coal  has  many  uses.  Production  of  anthracite  coal  is  now  about 
40  million  tons  per  year.  About  three-fourths  of  this,  output  goes  to 
the  retail  market  to  be  used  as  fuel  for  commercial  and  institutional 
buildings,  as  well  as  for  heating  homes.  The  remainder  of  the  yearly 
tonnage  is  used  by  a  variety  of  industrial  consumers. 

Annual  production  of  bituminous  coal  is  now  about  450  million 
tons  per  year.  The  major  use  of  this  coal  is  in  generating  electricity. 




About  one-half  of  our  nation's  total  electric  power  is  produced  by 
coal-fired  steam  plants.  In  all,  the  electric  utilities  consume  consider- 
ably more  than  100  million  tons  of  coal  each  year,  or  about  25  per- 
cent of  all  bituminous  coal  marketed.  The  production  of  electric 
power  is  expanding  rapidly  and  it  is  expected  that  there  will  be  an 
increase  in  the  amount  of  coal  used  for  this  purpose. 

The  steel  industry  is  the  second  most  important  customer  for 
bituminous  coal,  consuming  about  20  percent  of  all  coal  marketed. 
Most  of  this  enormous  quantity  of  coal  is  converted  into  coke  by 
roasting  in  coke  ovens.  Coke  is  used  to  some  extent  as  a  fuel,  but  its 
chief  use  by  far  is  as  a  reducing  agent  in  the  production  of  steel. 
Coke  robs  the  iron  ore,  an  oxide  of  iron,  of  its  oxygen  and  leaves  pig 
iron  which  is  later  converted  to  steel.  The  chemicals  recovered  in 
the  coking  process  (see  page  327)  are  the  source  of  many  plastics, 
dyes,  drugs,  medicines,  and  industrial  chemicals.  Today,  coal  chemi- 
cals are  also  produced  directly  from  the  "raw"  coal  without  any  at- 
tempt at  burning  it  or  changing  it  to  coke.  This  use  of  coal  as  yet 
constitutes  a  very  small  market,  but  it  may  become  increasingly  im- 
portant in  the  future.  Coal  is  also  destructively  distilled  in  the  manu- 
facture of  coal  gas. 

All  other  types  of  manufacturing  consume  another  15  to  20  per- 
cent of  the  bituminous  coal  marketed  annually.  A  similar  amount 
goes  to  the  retail  market  for  use  as  a  fuel  in  public  buildings,  com- 
mercial establishments,  and  homes.  In  addition  to  coal  in  its  conven- 
tional lump  form,  the  retail  market  each  year  uses  over  a  million 
tons  of  coal  in  the  form  of  fuel  briquets.  Fuel  briquets  are  pressed 
cubes,  cylinders,  or  ovoids  of  very  fine  bituminous  coal  (slack)  or  lig- 
nite held  together  by  a  burnable  binder. 

The  railroads,  once  the  coal  industry's  best  customer,  now  use 
considerably  less  than  ten  percent  of  the  total  annual  tonnage  as  the 

Consumers  Power  Company 

The  importance  of  coal  in 
generating  electricity  is  illus- 
trated by  this  coal  storage 
pile  at  a  modern  power 
station.  This  plant  burns 
about  675,000  tons  of  coal 
each  year. 


coal-burning  "iron  horse"  has  been  displaced  by  the  more  efficient 
diesel-electric  locomotive. 

Buying  coal  for  your  home.  The  most  important  consideration  is 
the  amount  of  heat  the  coal  will  produce.  Your  coal  dealer  usually 
knows  the  Btu  value  of  the  coal.  Such  information  may  also  be  ob- 
tained from  the  United  States  Bureau  of  Mines,  from  your  state  bu- 
reau of  mines,  or  from  the  company  that  mined  the  coal. 

If  you  know  the  number  of  Btu  in  a  pound  of  coal,  the  product  of 
this  number  multiplied  by  2000  is  the  number  of  Btu  per  ton.  The 
quotient  of  the  number  of  Btu  per  ton  divided  by  the  cost  per  ton  in 
dollars  is  the  number  of  Btu  per  dollar.  As  far  as  heat  value  alone  is 
concerned,  the  coal  with  the  greatest  number  of  Btu  per  dollar  is  the 
best  buy.  Expressed  as  a  formula,  this  method  of  comparing  coals 
may  be  written: 

Btu  per  Ib.  X  2000         _  ,  „ 

— - : — rr; =  Btu  per  dollar 

Cost  per  ton  in  dollars  r 

However,  as  anyone  who  has  ever  carried  out  "ashes"  knows,  the  ash 
content  of  coal  is  an  important  consideration  also.  One  coal  with  a 
slightly  higher  Btu-per-dollar  value  than  a  second  coal  actually  may 
not  be  as  desirable  as  the  second  coal,  if  its  ash  content  is  appreci- 
ably higher.  Moisture  content  and  dust  content  are  two  other  fac- 
tors that  must  be  considered.  Coal  dealers  may  control  the  dustiness 
of  coal  either  by  sprinkling  the  coal  with  water  or  by  spraying  the 
coal  with  special  oils.  From  the  point  of  view  of  the  consumer,  which 
of  these  methods  is  preferable?  Why? 

In  buying  coal,  it  is  advisable  to  buy  the  size  and  grade  recom- 
mended by  the  manufacturer  of  the  equipment  in  which  you  intend 
to  burn  it.  Buying  the  appropriate  size  and  grade  of  coal  will  enable 
you  to  get  more  Btu  out  of  the  coal  you  buy.  Most  coal-burning 
equipment  is  designed  to  operate  most  efficiently  with  coal  of  cer- 
tain size  and  grade. 

Because  of  seasonal  demands,  coal  is  usually  lower  in  cost  during 
summer  than  in  winter.  Since  coal  does  not  deteriorate  appreciably 
in  dry  storage,  for  example,  in  a  coal  bin,  it  is  good  policy  to  buy 
next  winter's  supply  in  the  summertime. 

Changing  energy  relationships.  The  sources  of  energy  upon  which 
our  civilization  depends  are,  as  we  know,  coal,  petroleum,  natural 
gas,  and  water  power.  Nuclear  energy  holds  great  promise,  but  its 
use  as  a  major  source  of  energy  is  still  in  the  future. 

Petroleum  and  natural  gas  constitute  but  a  tiny  part  of  our  na- 
tion's total  fuel  deposits.  Taken  together,  they  m^te  up  less  than 


two  percent  of  the  known  reserves  of  all  fuels.  Of  the  remaining  98 
percent,  oil  shale  forms  eight  percent  and  coal  forms  90  percent. 
More  than  92  percent  of  the  total  coal  reserve  is  bituminous  and 
lignite;  the  remainder  is  anthracite.  Although  the  experts  disagree 
on  the  exact  number  of  years  involved,  it  is  generally  accepted  that, 
at  our  present  rate  of  use,  our  reserves  of  high-quality  petroleum  and 
natural  gas  may  be  consumed  in  less  than  half  a  century,  while  our 
coal  reserves  are  adequate  for  several  hundred  years.  Oil  shale  is  an 
unexploited  resource  and  we  do  not  yet  know  the  extent  of  its  use- 

Because  of  coal's  abundance,  it  seems  likely  that  it  will  see  in- 
creased use  in  the  future  as  oil  and  gas  reserves  shrink.  Because  our 
society  is  completely  dependent  upon  a  plentiful  and  continuous  sup- 
ply of  energy,  proper  use  and  conservation  of  our  fuel  reserves  is  a 
problem  which  deserves  the  thoughtful  attention  of  every  citizen. 

The  energy  of  running  water,  while  limited  in  amount,  is  never- 
theless almost  limitless  in  duration.  For,  in  harnessing  such  energy, 
nothing  is  actually  used  up  that  will  not  occur  again.  The  water  cycle 
takes  care  of  this  problem  for  us.  However,  water  power  suffers  from 
a  serious  drawback  in  that  it  is  not  constant.  That  is,  during  certain 
seasons  of  the  year,  more  water  is  available  than  at  others.  Even 
dams  and  huge  reservoirs  have  not  yet  succeeded  in  solving  this  prob- 
lem completely,  but  they  have  made  tremendous  strides  in  doing  so, 
and  more  and  more  water  power  is  being  harnessed  as  the  years  go 
by.  However,  even  if  developed  to  the  fullest  extent,  our  water  re- 
sources could  not  provide  more  than  a  small  fraction  of  our  total 
annual  energy  requirements. 

The  utilization  of  hydroelectric  power  had  to  wait  for  science  to 
develop  methods  of  sending  electric  energy  over  long  distances,  and 
also  for  the  building  of  immense  dams  and  hydroelectric  plants. 
Recently,  the  capacity  of  existing  hydroelectric  power  plants  in  the 
United  States  totaled  25  million  horsepower,  while  the  undeveloped 
water  power  was  estimated  at  117  million  horsepower.  Engineers  be- 
lieve that,  long  before  we  can  no  longer  depend  upon  coal,  petro- 
leum, natural  gas,  and  the  harnessed  power  of  running  water,  nuclear 
energy  will  do  much  of  the  world's  work. 

Motor  fuels  from  coal.  Chemists,  faced  with  an  imminent  world- 
shortage  of  petroleum,  turned  to  the  development  of  methods  of 
synthesizing  essential  liquid  fuels.  One  of  the  most  widely  used  de- 
velopments has  been  the  changing  of  low-grade  coal  into  gasoline 
and  other  petroleum  products.  The  most  successful  results  were  ob- 
tained by  the  German  chemists,  Friedrich  Bergius  (ber'gi-obs) ,  Franz 



Fischer,  and  Hans  Tropsch.  In  1931,  Bergius  was  awarded  the  Nobel 
prize  in  chemistry. 

Bergius  found  that  coal  contained  half  as  much  hydrogen  as  petro- 
leum. By  doubling  the  amount  of  hydrogen  chemically  combined  in 
coal,  he  hoped  that  lie  could  obtain  a  product  that  would  resemble 
gasoline.  He  developed  a  process  now  called  the  Bergius  process,  in 
which  a  thick  paste  of  powdered  coal  mixed  with  heavy  oil  is  intro- 
duced into  steel  drums  and  heated  to  a  temperature  of  about  400°C. 
under  about  200  atmospheres  pressure.  Hydrogen  is  forced  into  the 
mixture,  which  also  contains  a  catalyst.  When  the  hydrogenation  is 
complete,  a  mixture  of  gasoline  and  heavier  fuel  oils  is  obtained. 
The  carbon  of  the  coal  combines  chemically  with  hydrogen,  form- 
ing certain  kinds  of  hydrocarbons,  and  the  successful  liquefaction 
of  coal  is  achieved. 

Assuming  the  composition  of  average  gasoline  to  be  represented 
by  the  formula  C7H1(J,  we  may  represent  the  hydrogenation  of  coal 
by  the  equation: 

7C+8H2-»C7H16  .-;"::;•:;  ' 

During  World  War  II,  the  Germans  demonstrated  that  synthetic 
gasoline  from  coal  can  be  used  for  the  same  purposes  as.  petroleum- 
derived  gasoline.  In  the  United  States,  following  the  war>  both  the 
federal  government  and  private  industry  conducted  research  on  the 
production  of  gasoline  from  coal  and  from  natural  gas.  It  was  found 
that  synthetic  gasoline  is  more  expensive  to  produce  than  gasoline 
from  petroleum,  but  would  be  commercially  practical  in  the  event 
of  shortages  of  the  latter. 

Apparatus  in  a  coal-industry 
research  laboratory  for  testing 
stoker  coal  and  thus  establish- 
ing standards  for  determining 
the  best  use  for  the  various 

The  testing  laboratory  of  a 
coal  preparation  plant.  The 
crucibles  contain  powdered 
coal  which  is  burned  in  the 
electric  furnace.  By  analysis 
of  the  ash,  the  composition 
of  the  coal  is  determined. 

National  Coal  Association 

Both  the  Bergius  and  Fischer-Tropsch  processes  were  tested.  The 
Fischer-Tropsch  process  depends  either  upon  the  reaction  of  steam 
and  coal,  or  oxygen  and  natural  gas,  to  form  synthesis  gas. 

CH-H20->CO  +  H2    or     2CH4  +  O2  -»  2CO  +  4H2 

synthesis  gas  synthesis  gas 

This  mixture  of  gases  is  then  converted  with  the  aid  of  a  catalyst 
into  gasoline. 

7  CO  +  15H2 >  C7H16  4-  7H2O 

• . catalyst    , 

synthesis  gas  synthetic 



Bragg,  William.  Concerning  the  Nature  of  Things.  Dover 
Publications,  New  York,  1954.  Lecture  IV  on  crystals  covers 
the  various  allotropic  forms  of  carbon.  A  classic. 

Storch,  H.  H.;  Lowry,  H.  H.;  Kiebler,  M.  W.,  Jr.;  How- 
ard, H.  C.;  Thiessen,  Gilbert;  and  Charmbury,  H.  B.  "Hy- 
drogenation  of  Coal."  Industrial  and  Engineering  Chemistry, 
April,  1944,  pp.  291-298.  A  round-table  discussion  of  problems 
in  this  process. 


1.  Destructive  distillation  is  the  heating  of  a  complex  or- 
ganic compound  in  a  chamber  from  which  air  is  excluded 
and  condensing  the  vapors  formed. 



2.  Fuel  value  of  coal  is  measured  in  calories  or  British 
thermal  units,  Btu.  One  Btu  is  the  amount  of  heat  required 
to  raise  1  Ib.  of  water  1°F.  A  calory  is  the  amount  of  heat  re- 
quired to  raise  1  g.  of  water  1°C.  One  Btu  is  equivalent  to 
252  calories. 

3.  Organic  chemistry  is  the  chemistry  of  carbon  compounds. 
At  present  more  than  650,000  compounds  of  carbon  are  known. 
Synthetic  chemistry  reaches  its  greatest  development  in  this 
branch  of  chemistry  and  has  already  produced  dyes,  drugs,  and 
solvents  many  of  which  are  better  than  natural  substances. 

4.  Today,  man's  chief  sources  of  energy  are  coal,  petroleum 
and  natural  gas,  and  water  power.  In  the  years  just  ahead, 
nuclear  energy  is  likely  to  assume  greater  importance. 


Group  A 

1.  C  may  act  either  as  a  metal  or  as  a  nonmetal.  (a)  Why? 
(b)  Why  is  C  not  active  at  ordinary  temperatures? 

2.  (a)  What  are  the  three  allotropic  forms  of  C?  (b)  Name 
the  allotropic  forms  of  two  other  elements. 

3.  (a)  What  chemical  property  is  common  to  all  allotropic 
forms  of  C?  (b)  What  product  depends  on  this  property? 

4.  What  accounts  for  the  extreme  hardness  of  diamond? 

5.  How  did  the  action  of  extreme  heat  on  diamond  aid 
Lavoisier  in  securing  acceptance  of  his  theory  of  burning? 

6.  (a)  What  are  two  possible  products  of  the  oxidation 
of  C?  (b)  Write  equations  to  illustrate. 

7.  (a)  What  are  the  three  chief  uses  of  diamond?  (b)  Upon 
what  properties  do  these  depend? 

8.  What  evidence  supports  the  accepted  theory  of  diamond 
formation?  ft 

9.    (a)  What   are   the   chief   uses  of  graphite?    (b)  Upon 
what  properties  do  these  depend? 

10.  Describe  briefly  the  commercial  preparation  of  graphite. 

11.  (a)  What  is  an  electric  furnace?  (b)  Name  three  kinds. 

12.  Make  a  diagram  of  one  kind  of  electric  furnace  and  ex- 
plain how  it  produces  high  temperatures. 

t .  .- 

13.  Make  a  table  of  four  kinds  of  amorphous  C,  giving  the 


method  of  manufacture,  composition,  and  chief  uses  of  each. 

14.  (a)  What  is  destructive  distillation?    (b)  Describe  the 
making  of  coke,  and  (c)  name  six  byproducts  of  the  process. 

15.  Activated  charcoal  is  used  in  gas  masks.  Why? 

16.  Below  are  two  lists  of  items.  Match  each  item  in  the  list 
on  the  left  with  the  correct  item  from  the  list  on  the  right. 

1)  watch  bearings  a)  activated  charcoal 

2)  Moissan  b)  graphite 

3)  peat  c)  printer's  ink 

4)  Bergius  d)  diamond  crystal 

5)  lampblack  e)  coke 

6)  manufacture  of  iron  f)  artificial  diamonds 

7)  lead  pencils  g)  bacteria 

8)  adsorption  h)  hydrogenation 

17.  What  weight  of  coke  containing  80  percent  C  is  needed 
to  reduce  5  Ib.  of  CuO? 

T    f  ,  T  , 


18.  Assume  the  formula  of  gasoline  to  be  C7H16.  If  500  ml. 
of  water  vapor  are  formed  during  the  combustion  of  a  certain 
quantity  of  gasoline,  what  volume  of  CO2  will  be  formed? 

19.  What  volume  of  CO2  will  be  formed  during  the  com- 
bustion oi;  2000  Ib.  of  coal  containing  80  percent  C?   (Note: 
the  ounce-molecular  volume  of  any  gas  is  22.4  cu.  ft.) 

20.  According  to  available  evidence,  how  were  coal  de- 
posits formed? 

21.  Compare   anthracite  and  bituminous  coal  for  use   in 
household  heating. 

22.  What  factors  should  be  considered  in  buying  coal? 

23.  How  may  the  heat  value  of  one  coal  be  compared  with 
the  heat  value  of  another? 

24.  Complete  and  balance  the  following  equations: 

a)  C  +  H2       ->  c) 

b)  Fe2O3  +  C  ->  d)  SiO2  +  C  -> 

25.  What  is  meant  by  saying  that  a  certain  sample  of  coal 
liberates  14,000  Btu? 

26.  What  is  the  major  use  of  anthracite  coal?  What  two  in- 
dustries are  the  major  consumers  of  bituminous  coal? 

„  .  I      . 

27.  Explain  the  Bergius  process  for  the  hydrogenation  of 


coal.  Why  did  he  experiment  to  achieve  this  result? 

28.  With   the   aid   of  equations   explain   the   synthesis   of 
gasoline  by  the  Fischer-Tropsch  process. 

29.  (a)  What  are  three  physical  properties  possessed  by  all 
allotropic  forms  of  C?   (b)  In  what  physical  property  do  they 
differ  radically? 

30.  Explain  why  the  term  lead  pencil  is  a  misnomer. 

31.  Give  two  reasons  why  graphite  is  used  in  making  stove 

Group  B 

32.  2  g.  of  coal  raised  the  temperature  of  2000  ml.  of  water 
6°C.  Determine  the  fuel  value  of  this  coal  in  calories. 

33.  Name  four  reducing  agents  and  give  an  example  of  the 
reducing  power  of  each. 

34.  How  is  an  imitation  diamond  distinguished  from  the 

35.  Wooden  posts  are  sometimes  charred  before  being  placed 
in  the  ground.  Explain. 

36.  Charcoal  tablets  are  sometimes  used  in   the  treatment 
of  certain  stomach  disorders.  Explain. 

37.  What  produces  chemical  changes  in  an  electric  furnace: 
the  heat  or  the  electric  current?  Explain. 

38.  How  does  a  large  percentage  of  ash  or  S  in  coal  affect 
its  quality  as  a  fuel? 

39.  (a)  Why  is   Woehler  sometimes  called   the   father  of 
organic  chemistry?  (b)  What  is  the  difference  between  organic 
and  inorganic  compounds? 

40.  How  would  you  identify  three  black  powders  in  a  mix- 
ture of  graphite,  manganese  dioxide,  and  copper  oxide? 


1.  Place  a  small  amount  ot  sugar  on  a  hot  stove  or  hot  plate 
and  report  what  happens,  indicating  as  many  changes  as  you 

2.  Make  a  small  working  model  of  an  electric  lurnace  used 
in  the  production  of  graphite. 

3.  Large  users  of  coal  purchase  coal  according  to  definite 
specifications  as  to  the  percentage  of  C,  volatile  material,  ash 
and  sulfur  content,  and  fuel  value.  Does  your  father  buy  coal 
according  to  definite  specifications?  Why  do  most  domestic 
consumers   not   bother   about  specifications   in   buying  coal? 
Is  this  a  sensible  state  of  affairs?  Explain. 



GAS    OF    LIFE    AND    DECAY 

.  .  .  In  /757  Joseph  Black  discov- 
ered that  carbonic  acid  gas  could  not 
be  breathed  by  animals,  and  had  a 
poisonous  effect  on  them.  Sparrows 
introduced  into  an  atmosphere  of 
the  gas  died  in  10  or  n  seconds,  but 
if  their  nostrils  were  stopped  with 
lard,  their  death  took  place  only  at 
the  end  of  3  or  4  minutes.  Robert 
Routledge  in  A  Popular  History  of 
Science,  1881 

A  Scottish  physician  discovers  carbon  dioxide.  Joseph  Black,  a 
physician,  was  very  much  interested  in  a  lively  discussion  between 
two  professors  of  medicine.  One  professor  believed  that  limewater 
made  from  limestone  was  a  more  effective  medicine  than  limewater 
made  from  oyster  shells.  Black,  to  settle  this  controversy  in  true  sci- 
entific manner,  undertook  a  thorough  study  of  several  carbonates. 
In  1754,  while  heating  one  of  these  carbonates  (magnesium  carbo- 
nate) ,  he  obtained  pure  carbon  dioxide  gas  for  the  first  time.  This 
was  20  years  before  the  isolation  of  pure  oxygen.  Black  called  car- 
bon dioxide  "fixed  air." 

MgCO3  -»  MgO  +  CO2  T 

The  atmosphere  is  a  vast  storehouse  of  this  gas.  Immense  quanti- 
ties also  occur  dissolved  in  surface  and  ground  water.  Large  quan- 
tities escape  from  volcanoes  and  crevices  in  the  earth;  it  is  also 
ejected  in  tremendous  volumes  from  wells  that  produce  petroleum 
or  natural  gas.  In  fact,  enough  carbon  dioxide  could  be  obtained 
from  oil  wells  in  the  United  States  alone  to  supply  the  world.  Carbon 
dioxide  is  formed  when  carbon  or  its  com  pounds  are  burned.  Be- 
cause all  common  fuels  contain  carbon  either  free  or  in  compounds, 



«..   ~  TH.  c 
oxygen  cycle. 

green  plants  in  sunlight 

carbon  dioxide  is  produced  whenever  such  fuels  are  burned.  Finally, 
tremendous  volumes  of  carbon  dioxide  are  locked  up  chemically  in 
the  great  masses  of  metallic  carbonates  and  bicarbonates  found  in 
the  earth's  crust. 

Why  carbon  dioxide  is  necessary  for  life.  Life,  as  we  know  it  today, 
would  be  impossible  without  this  percentage  of  carbon  dioxide  gas 
in  the  air,  small  though  it  is  (0.04  percent) .  For  without  it,  plants 
even  with  the  aid  of  sunlight  could  not  manufacture  starch.  As  a 
result,  our  food  supply  would  diminish  and  finally  vanish  altogether. 
Since  coal  was  formed  in  prehistoric  days  by  the  destructive  distilla- 
tion of  plants,  even  this  great  source  of  power  would  never  have  been 
formed  if  carbon  dioxide  had  not  been  present  in  air. 

On  the  other  hand,  a  much  larger  concentration  of  carbon  dioxide 
in  the  atmosphere  would  be  fatal,  for  it  would  dilute  the  oxygen  and 
choke  the  breath  of  life  out  of  living  things.  The  delicately  balanced 
ratio  of  oxygen  to  carbon  dioxide  in  the  air  may  be  better  appreciated 
after  a  study  of  the  carbon  dioxide-oxygen  cycle. 

The  carbon  dioxide-oxygen  cycle.  Carbon  dioxide  is  constantly 
being  added  to  the  air  by  the  breathing  of  animals.  It  is  made  in  the 
body  tissues  by  the  oxidation  of  carbon  compounds  in  these  tissues. 
Some  of  the  oxygen  inhaled  in  air  is  changed  to  carbon  dioxide, 
which  passes  out  of  the  body  through  the  lungs.  Every  breath  of  air 
exhaled  contains  100  times  as  much  carbon  dioxide  as  was  inhaled. 
At  the  same  time,  the  amount  of  oxygen  decreases  from  about  2 1  per- 
cent in  the  air  inhaled  to  about  16  percent  in  the  air  exhaled. 

Every  ton  of  coal  and  liquid  fuel  burned  sends  about  four  tons  of 
carbon  dioxide  into  the  air,  diminishing,  at  the  same  time,  the  oxy- 
gen content  of  the  air.  During  the  decay  and  fermentation  of  or- 
ganic material,  immense  volumes  of  carbon  dioxide  are  given  off 
into  the  air. 

You  might  suspect  that,  after  hundreds  of  thousands  of  years, 
these  three  processes  — -  breathing,  burning,  and  decay  —  would  have 
filled  the  air  with  enough  carbon  dioxide  to  destroy  life.  But  com- 
pensating devices  offset  this  increase  in  carbon  dioxide. 




All  the  waters  of  the  earth  are  continually  dissolving  carbon  di- 
oxide, thus  removing  it  i'rom  air.  In  the  formation  of  carbonates  and 
bicarbonates  in  nature,  great  quantities  of  free  carbon  dioxide  of  the 
air  are  used  up  in  the  chemical  weathering  of  rocks.  Most  important 
of  all,  during  daylight  hours  all  green  plants  are  absorbing  carbon 
dioxide,  converting  it  into  a  form  of  sugar  (fruit  sugar)  and  into 
starch,  and  returning  free  oxygen  to  the  air.  The  process  that  is 
responsible  for  this  change,  photosynthesis,  has  been  described  by 
some  writers  as  the  most  important  chemical  process  in  the  world, 
for  if  it  should  cease  man  and  other  animals  would  have  no  food. 

In  higher  plants,  photosynthesis  takes  place  only  in  parts  that  con- 
tain chlorophyll,  a  green  coloring  matter  that  acts  as  a  catalyst.  The 
final  equation  for  the  formation  of  starch  is: 

6CO2  +  5H2O  — »  6O2  -f  C6H,oO5  (starch) 

In  Chapter  36  photosynthesis  in  relation  to  food-building  is  dis- 
cussed, but  at  present  we  are  concerned  particularly  with  its  effect 
upon  the  composition  of  air  and  the  way  in  which  it  helps  to  main- 
tain the  balanced  ratio  of  carbon  dioxide  to  oxygen.  This  balance, 
as  it.  relates  to  plant  and  animal  life,  is  indicated  by  Fig.  74. 

How  carbon  dioxide  is  prepared.  When  an  acid  is  added  to  a 
carbonate  or  bicarbonate,  carbon  dioxide  is  liberated.  In  the  labora- 
tory, calcium  carbonate,  in  the  form  of  chips  of  marble  or  limestone, 
is  treated  with  hydrochloric  acid,  and  the  CCX  is  collected  by  water 

In  these  tanks,  chlorophyll,  the  catalytic  agent  in  photosynthesis, 
is  extracted  from  alfalfa  leaves. 

American  Chlorophyll  Division  Strong,   Cobb   and  Company,   Inc. 


displacement,  as  shown  in  Fig.  75  below.  This  gas  may  also  be  col- 
lected by  air  displacement,  for  CO2  is  heavier  than  air. 

CaCO3  +  2HC1  -»  CaCl2  +  H2O  +  CO2 1 

Most  of  the  carbon  dioxide  used  commercially  is  obtained  as  a 
byproduct  from  coke  furnaces,  gas  wells,  and  fermentation  processes. 
The  heating  of  limestone  to  make  quicklime 

CaCO3  -»  CaO  +  CO2 

also  furnishes  carbon  dioxide  for  industrial  use.  Carbon  dioxide 
for  industrial  use  is  separated  thoroughly  from  all  impurities  before 
being  liquefied  and  stored  in  steel  cylinders  under  pressure.  It  is 
known  commercially  as  carbonic  acid  gas. 

Physical  properties  of  carbon  dioxide.  Carbon  dioxide  is  a  color- 
less, odorless  gas;  it  is  1|  times  as  heavy  as  air,  and  is  soluble  in  wa- 
ter volume  for  volume  (at  room  temperature  and  pressure) .  At 
higher  pressures  and  lower  temperatures,  it  dissolves  in  water  in 
much  greater  volumes. 

At  ordinary  temperatures,  a  pressure  of  52  atmospheres  is  suffi- 
cient to  condense  carbon  dioxide  to  a  colorless  liquid.  When  the 
pressure  over  this  colorless  liquid  is  decreased,  it  evaporates  rapidly. 
During  this  process,  evaporation  of  part  of  the  liquid  withdraws  heat 
so  rapidly  from  the  remaining  liquid  that  a  great  amount  of  the 
liquid  carbon  dioxide  is  changed  to  solid  carbon  dioxide,  called 
carbon  dioxide  snow  or  dry  ice. 

Chemical  properties  of  carbon  dioxide.  Carbon  dioxide  does  not 
burn.  It  will  support  the  combustion  of  only  very  active  substances, 
such  as  sodium.  Because  it  is  heavier  than  air,  it  can  be  poured  like 
a  liquid.  When  a  bottle  of  the  gas  is  poured  over  a  lighted  candle, 
the  flame  is  extinguished.  Why? 

The  chief  chemical  property  of  carbon  dioxide  is  its  ability  to 
combine  with  water,  forming  carbonic  acid.  When  carbon  dioxide 
gas  is  bubbled  through  water,  part  of  it  combines  chemically  and 


Fig.  75.  Laboratory  prepara- 
tion of  carbon  dioxide.  What 
other  gat  may  be  collected 


Fig.  76.  A  cubical  diagram 
of  the  carbon  dioxide  mole- 




forms  carbonic  acid;  the  rest  merely  dissolves  in  the  water.  Carbon 
dioxide  is,  therefore,  the  anhydride  of  carbonic  acid. 

H2O  +  CO2  ->  H2CO3  (carbonic  acid) 

Electron  structure  of  carbon  dioxide.  Carbon  dioxide  is  a  non- 
polar,  covalent  compound  whose  structure  may  be  represented  as 
shown  in  Fig.  76.  Notice  that  each  oxygen  atom  shares  two  pairs  of 
electrons  with  the  carbon  atom.  The  carbon  atom  shares  four  pairs  of 
electrons  with  the  two  oxygen  atoms.  Since  the  valence  of  an  ele- 
ment in  a  nonpolar  compound  is  equal  to  the  number  of  shared 
pairs  of  electrons,  the  valence  of  oxygen  is  two  and  the  valence  of 
carbon  is  four. 

Carbonic  acid.  Carbonic  acid  closely  resembles  sulfurous  acid.  It 
is  unstable,  decomposing,  when  warmed,  into  water  and  its  anhy- 
dride, CO2.  Thus  the  reaction  for  the  formation  of  carbonic  acid  is 

H2C03  *±  H20  +  C02 

Because  of  the  highly  unstable  nature  of  carbonic  acid,  it  has  never 
been  isolated  in  the  pure  state. 

Like  sulfurous  acid,  carbonic  acid  is  a  weak  acid,  dissociating  only 
slightly,  forming  few  hydrogen  ions.  Carbonated  waters  containing 
dissolved  carbon  dioxide  are  so  weakly  acid  that  they  may  be  con- 
sumed in  large  quantities  without  ill  effect.  Although  carbonic  acid 
turns  blue  litmus  paper  pink,  its  sour  taste  is  scarcely  noticeable. 

Some  of  the  many  uses  of  carbon  dioxide.  The  chief  uses  of  carbon 
dioxide  are:  (1)  in  the  manufacture  of  effervescent  beverages,  (2)  as 
dry  ice  for  refrigeration,  (3)  in  the  leavening  of  bread,  (4)  in  extin- 
guishing fires,  (5)  in  the  manufacture  of  washing  and  baking  soda, 
and  (6)  in  the  modern  synthesis  of  urea,  (NH2)  2CO,  a  white  crystal- 
line substance  used  in  the  impregnation  of  wood,  in  the  treatment 
of  wounds  (under  the  name  carbamide) ,  and  in  making  a  valuable 
fertilizer,  "Uramon": 

2NH3  +  C02  ->  H20  +  (NH2)2CO 

The  normal  human  adult  eliminates  about  30  grams  of  urea  as  a 
waste  product  daily  in  liquid  excretions. 



Rubber  life  rafts  and  life  belts  are  inflated  by  means  of  carbon 
dioxide.  The  carbon  dioxide  so  used  is  stored  under  pressure  in  small 
steel  bulbs  not  unlike  the  carbon  dioxide  cartridges  used  on  siphons 
in  making  charged  water. 

Another  use  of  carbon  dioxide  is  in  making  car  bo  gen,  a  mixture 
of  95  percent  oxygen  and  five  percent  carbon  dioxide,  administered  to 
victims  of  gas  poisoning  and  pneumonia  to  induce  more  rapid  res- 
piration. Carbon  dioxide  in  the  blood  appears  to  stimulate  the  re- 
spiratory nerve  centers  that  control  breathing.  The  heavy  breathing 
of  a  runner,  for  example,  is  caused  by  the  large  quantity  of  carbon 
dioxide  produced  in  his  body  during  violent  exercise.  Adding  more 
carbon  dioxide  to  the  air  inhaled,  therefore,  causes  the  more  rapid 
breathing  needed  by  a  patient  under  treatment. 

How  carbon  dioxide  is  used  in  effervescent  beverages.  Priestley 
discovered  the  pleasant  taste  of  water  containing  dissolved  carbon 
dioxide.  Before  the  Royal  Society  of  England  he  prepared  "a  glass 
of  exceedingly  pleasant  sparkling  water  which  could  hardly  be  dis- 
tinguished from  Seltzer  water,"  and  received  the  Society's  gold  medal 
for  his  discovery.  This  was  the  first  great  and  unforgettable  triumph 
of  this  amateur  chemist. 

Years  later,  when  he  was  forced  to  come  to  America,  Priestley  in- 
terested Dr.  Physick,  of  Philadelphia,  in  this  beverage.  In  1807, 
Dr.  Physick  had  a  chemist  prepare  carbonated  water  with  a  little 
fruit  juice  for  his  patients.  This  was  the  beginning  of  the  soft  drink 
industry  in  America,  which  uses  annually  more  than  200  million 
pounds  of  carbon  dioxide. 

Soda  water,  so  named  because  sodium  carbonate  was  then  used 
in  preparing  the  carbon  dioxide,  is  now  prepared  by  forcing  carbon 
dioxide  gas  into  cold  water  at  high  pressures.  When  the  pressure  is 
released,  the  excess  carbon  dioxide  gas  is  liberated,  causing  the  bub- 
bling, or  effervescing.  Carbon  dioxide  is  supplied  to  soda  fountains 
in  liquid  form  in  steel  cylinders. 

Dry  ice.  White,  solid  carbon  dioxide  has  come  into  wide  use  for 
the  refrigeration  of  foods  especially  perishable  products  in  transit. 
It  has  three  advantages  over  ice:  (1)  the  temperature  of  dry  ice, 
—  78°C.,  is  much  lower  than  that  of  ice;  (2)  dry  ic$  does  not  melt 
into  a  liquid,  but  changes  directly  into  a  gas;  and  (3)  in  changing 
from  a  solid  to  a  gas,  dry  ice  absorbs  three  times  as  much  heat  ^s  ice 
does  when  it  melts. 

Some  dry  ice  is  used  in  the  cooling,  and  hardening  of  rivets  made 
of  aluminum  alloys,  and  in  shrink-fitting.  This  is |a  process  by  which 
a  metal  fitting  of  correct  size  is  expanded  by  hel$i|kg  until  it  can 

Blocks  of  carbon  dioxide  dry  ice, 
direct  from  the  hydraulic  presses, 
are  cut  into  cubes  before  being 
wrapped  for  shipment. 

•ill   Corpnr 

be  placed  over  or  around  the  base  to  which  it  is  to  be  attached.  The 
fitting  is  then  shrunk  by  cooling  until  it  adheres  to  the  base  with 
great  pressure.  Dry  ice  is  also  used  in  the  low-temperature  drying  of 
many  different  kinds  of  biological  materials,  in  the  preservation  and 
shipment  of  blood  plasma,  and  in  the  quick-free/ ing-  of  many  dif- 
ferent kinds  of  foods. 

Baking  powders  liberate  carbon  dioxide  lor  leavening.  Bread 
that  is  not  porous  is  hard,  unpalatable,  and  somewhat  indigestible. 
Leavening  bread  makes  it  porous,  light,  and  more  easily  digested. 
Although  bread  may  be  leavened  by  beating  air  into  the  dough,  the 
method  more  generally  used  is  to  liberate  large  volumes  of  carbon 
dioxide  gas  in  the  dough  by  chemical  action.  Carbon  dioxide  is  an 
ideal  gas  for  leavening  because  it  is  colorless,  odorless,  nonpoisonous, 
and  easily  and  inexpensively  prepared. 

It  is  generally  produced  by  baking  powders,  which  are  mixtures 
of  two  white  powders,  one  of  which  is  sodium  bicarbonate,  or  baking 
soda,  and  the  other  a  substance  such  as  monocalcium  phosphate, 
Ca  (H,,PO4)  .„  cream  of  tartar,  KHC4H4O,.,  or  sodium  aluminum  sul- 
fate,  NaAl  (SO4)  2.  While  dry,  the  two  powders  do  not  react,  since 
most  inorganic  reactions  take  place  between  ions.  However,  when 
water  is  added,  the  powders  dissolve,  dissociate,  and  an  ionic  reaction 
takes  place  with  the  liberation  of  carbon  dioxide  gas,  causing  the 
reaction  to  go  to  completion.  One  of  these  reactions  may  be  repre- 
sented as  follows: 

Baking  soda  +  cream  of  tartar  — >  Rochelle  salt 

NaHCO3  +  KHC4O6  ->  KNaC,H4O6  +  H2O  +  CO2  \ 

Sodium  bicarbonate  +  potassium  acid  — >  potassium  sodium 
tartrate  tartrate 

Baking  powders  differ  mainly  in  their  speed  of  action.  They  also 
contain  about  15  percent  starch  or  flour  as  a  filler  to  keep  the  salts 
dry  and  thus  prevent  them  from  reacting  before  being  added  to  the 



concentrated  H2SO4 

Fig.   77.    Cross   section   of  a  portable 

carbon  dioxide  fire  extinguisher.  concentrated 

solution  of  NaHCO3 

copper  tank  -» 

Some  housewives  prefer  to  use  their  own  ingredients  in  leavening. 
For  example,  in  making  sour-milk  biscuits,  which  many  people  pre- 
fer to  baking-powder  biscuits,  the  housewife  uses  baking  soda  and 
sour  milk,  which  contains  lactic  acid. 

Carbon  dioxide  for  leavening  is  produced  also  by  yeast,  one  of 
the  oldest  of  leavening  agents.  Yeast  consists  of  living  plant  cells, 
which  produce  zymase.  Zymase,  which  is  a  mixture  of  several  enzymes, 
acts  catalytically  on  the  sugar  present  in  the  dough  and  breaks  it 
down  into  alcohol  and  carbon  dioxide  (see  page  548) .  What  advan- 
tage has  yeast  over  commercial  baking  powders? 

How  carbon  dioxide  is  used  in  firefighting.  Carbon  dioxide  gas  is 
an  ideal  firefighter.  It  does  not  support  combustion,  is  heavier  than 
air,  and  can  be  quickly  and  cheaply  liberated  in  large  volumes. 

One  kind  of  common  portable  fire  extinguisher  consists  of  a  copper 
tank  partly  filled  with  a  concentrated  solution  of  sodium  bicarbonate. 
Resting  on  a  shelf  inside  the  top  of  the  tank  is  a  bottle  of  concen- 
trated sulfuric  acid  covered  with  a  loose  lead  stopper.  When  the  tank 
is  inverted,  the  sulfuric  acid  pours  out  of  the  bottle,  and  reacts  with 
the  sodium  bicarbonate,  liberating  carbon  dioxide  gas,  which  carries 
out  of  the  nozzle  with  it  a  fine  spray  of  water  and  some  sodium  sul- 
fate.  The  equation  is: 

2NaHCO3  +  H2SO4  ->  Na2SO4  +  2H2O  +  2CO2 1 

To  put  out  stubborn  fires,  such  as  oil  conflagrations  which  water 
cannot  extinguish,  "Foamite-Firefoam"  mixture  is  used.  In  this,  an 
aluminum  sulfate  solution  takes  the  place  of  the  sulfuric  acid  in  the 
ordinary  type  of  fire  extinguisher.  This  aluminum  salt  combines 
with  water  by  a  process  called  hydrolysis,  forming  gelatinous  alumi- 
num hydroxide.  The  carbon  dioxide  liberated  is  held  fast  in  large, 
tough  bubbles  by  the  aluminum  hydroxide  as  well  as  by  a  sticky 
extract  of  licorice.  Thus,  the  mixture  spreads  over  the  fire  a  layer 
of  large  bubbles  of  carbon  dioxide.  These  smother  the  fire  by  keep- 
ing out  the  oxygen  of  the  air.  The  equation  is: 

A12(SO4)3  +  6NaHCO3  ->  2A1(OH)3  +  6CO2  +  3Na2SO4 


Fighting  fire  with  liquid  carbon  dioxide.  Another  form  of  fire- 
fighting  apparatus  used  very  widely  depends  on  liquid  carbon  diox- 
ide under  pressure  in  steel  containers.  The  liquid  carbon  diox- 
ide may  be  in  portable  cylinders  or  in  built-in  systems  (such  as  in 
airplane  engines) .  When  the  valve  is  opened,  the  liquid  carbon  di- 
oxide vapori/es  and  rushes  very  rapidly  out  of  the  tube.  This  sud- 
den evaporation  of  part  of  the  carbon  dioxide  cools  the  remaining 
liquid  to  a  white  solid,  and  this  carbon  dioxide  snow,  played  over 
the  fire,  quickly  puts  it  out.  The  fact  that  carbon  dioxide  is  able  to 
penetrate  obstructions  without  damaging  equipment  makes  it  one 
of  the  most  rapid  and  efficient  of  fircfighting  substances. 

Many  ships,  including  the  Queen  Mary,  are  equipped  to  use  car- 
bon dioxide  for  fighting  fires  in  the  various  storage  compartments 
throughout  the  ship.  In  a  recent  year,  more  than  500  million  pounds 
of  liquid  carbon  dioxide  and  carbon  dioxide  gas  were  produced  in 
this  country. 

Salts  of  carbonic  acid.  Because  carbonic  acid  contains  two  replace- 
able hydrogen  atoms,  it  is  a  dibasic  acid.  As  you  know,  one  or  both 
of  these  atoms  may  be  replaced  by  a  metal.  Hence  carbonic  acid 
forms  two  series  of  salts,  carbonates  and  bicarbonates.  Thus  carbonic 
acid  reacts  with  sodium  hydroxide,  forming  sodium  carbonate  or 
sodium  bicarbonate,  depending  upon  the  conditions. 

H2CO3  +  2NaOH  ->  Na2CO3  (sodium  carbonate)  +  2H2O 
H2CO3  +  NaOH  ->  NaHCO3  (sodium  bicarbonate)  +  H2O 

With  the  exception  of  sodium  carbonate,  potassium  carbonate,  and 
ammonium  carbonate,  all  carbonates  are  insoluble  in  water.  When  a 

Demonstrating  the  effectiveness 
of  a  carbon  dioxide  extinguisher 
against  a  fire  in  electrical  equip- 



carbonate  is  heated,  carbon  dioxide  is  liberated  and  the  oxide  of 
the  metal  remains. 

CaCO3  -»  CaO  +  CO2  f 

The  carbonates  of  sodium  and  calcium  are  discussed  in  Chapters  30 
and  31. 

How  to  test  for  a  carbonate  or  bicarbonate.  An  acid  liberates 
carbon  dioxide  gas  from  either  a  carbonate  or  a  bicarbonate.  If  this 
gas  is  then  passed  into  limewater,  Ca  (OH)  2,  a  white  precipitate,  cal- 
cium carbonate,  CaCO3,  forms. 

Ca(OH)2  +  C02  ->  CaC03 1  +  H2O 

On  bubbling  more  carbon  dioxide  gas  through  the  mixture,  the 
white  precipitate  disappears  as  a  result  of  the  solubility  of  calcium 
bicarbonate  which  is  formed. 

CaCOg  +  H2O  +  CO2  ->  Ca(HCO3)2  (calcium  bicarbonate) 


Armstrong,  George  B.  "Dry  Ice."  Chemistry,  Feb.  1951, 
pp.  24-31.  An  excellent,  illustrated  article  on  the  history, 
properties,  manufacture  of  and  uses  of  solid  carbon  dioxide. 

U.S.  Department  of  Agriculture,  1945.  Safe  Use  and  Storage 
of  Gasoline  and  Kerosene.  Farmers'  Bulletin  No.  1678.  Supt.  of 
Documents,  Washington,  D.C.  10^.  Some  excellent  advice  on 
putting  out  gasoline  and  kerosene  fires. 




and  . 



Group  A 

What  are  the  chief  sources  of  CO2? 

How  and  by  whom  was  CO2  discovered? 

In  what  way  does  life  depend  upon  CO2? 

Copy  and  complete  the  following  statements.  Do  not 

in  this  book.  CO2  is  being  added  to  the  air  constantly 

. ,  .  .  . ,  and  ....  CO2  is  removed  from  the  air  by  .  .  . ,  .  .  . , 

Make  a  diagram  representing  the  carbon  dioxide-oxygen 
By  what  three  commercial  methods  is  CO2  produced? 


7.  Write  the  equation  for  one  commercial  method  of  pro- 
ducing CO2. 

8.  Make  a  labeled  diagram  for  the  laboratory  preparation 
of  CO2. 

9.  (a)  Write   the   chemical   equation   for   the   laboratory 
preparation  of  CO2.    (b)  Why  does  this  reaction  go  to  com- 

10.  (a)  Can  carbon  dioxide  be  collected  by  air  displace- 
ment? (b)  Explain. 

11.  What  are  the  physical  properties  of  CO2? 

12.  What  are  three  chemical  properties  of  CO2? 

13.  (a)  Write  the  reversible  reaction  of  carbon  dioxide  and 
water,   (b)  How  could  you  make  large  quantities  of  CO2  com- 
bine with  H2O?  (c)  The  behavior  of  CO2  with  water  places  it 
in  what  class  of  compounds? 

14.  What  other  acid  closely  resembles  H2CO3? 

15.  (a)  What  are  six  uses  of  CO2?    (b)  Opposite  each  use 
state  the  property  or  properties  on  which  that  use  depends. 

16.  Describe  briefly  the  beginning  of  the  soft-drink  industry 
in  America. 

17.  (a)  How  is  dry  ice  made?   (b)  Dry  ice  has  what  advan- 
tages over  ice?  t 


18.  What  is  the  general  composition  of  baking  powder? 

19.  What  is  the  function  of  each  ingredient  in  one  type  of 
baking  powder? 

20.  Write  a  balanced  equation  representing  the  liberation 
of  carbon  dioxide  from  one  type  of  baking  powder. 

21.  (a)  Describe  the  chemical  action  of  some  other  type  of 
leavening  agent,    (b)  What  advantage  has  it  over  commercial 
baking  powders? 

22.  What  are  four  reasons  that  CO2  is  an  ideal  leavening 

23.  A  manufacturer  makes  enough  baking  powder  to  pro- 
duce 15  tons  of  CO2.  What  weight  of  NaHCO3  does  he  use? 

24.  A  sample  of  baking  powder  liberated  200  ml.  of  CO2. 
Considering  that  the  reaction  went  to  completion,  what  weight 
of  NaHCO3  was  used  if  the  baking  powder  was  of  the  cream- 
of- tartar  type? 



25.  What  volume  of  air  would  be  needed  to  burn  completely 
enough  coke  to  make  10,000  cu.  ft.  of  CO2? 

26.  Why  do  baking  powder  reactions  go  to  completion? 

27.  Describe  with  diagram  and  equation  the  principle  of 
one  type  of  portable  fire  extinguisher. 

28.  (a)   H2O  is  not  a  good  agent  for  putting  out  a  fire  of 
burning  oil.  Why?   (b)  Why  is  CO2  from  the  ordinary  type  of 
fire  extinguisher  ineffective?  (c)  What  principle  is  used  in  the 
"Foamite-Firefoam"  method? 

29.  (a)  H2CO3    has  two   series   of   salts.    What    are   they? 
(b)  Write  two  equations  to  illustrate. 

30.  What  is  the  action  of  heat  on  carbonates? 

31.  (a)  Describe  a  complete  test  for  a  carbonate,  (b)  Write 
the  equation  for  this  chemical  test. 

32.  When  sour  milk  is  used  in  cooking,  NaHCO3  is  used 
instead  of,  or  in  addition  to,  baking  powder.  Explain. 

33.  How  can  CO2  be  removed  from  a  sample  of  air? 

34.  Can  pure  GO2  alone  be  used  in  fighting  fire?  Explain. 

Group  B 

35.  (a)  Why  does  CO2  collect  in  wells  and  caves?  (b)  What 
is  its  source? 

36.  Does  the  effervescence  of  a  solution  when  an  acid  is 
added  prove  the  liberation  of  CO2?  Explain. 

37.  250  ml.  of  a  gas  weigh  0.49  g.  This  gas  contains  27.3 
percent  carbon  and  72.7  percent  oxygen.  Find  the  formula  of 
this  compound. 

38.  Most  chemical  reactions  take  place  between  ions.  How- 
ever, CaCO3  is  insoluble  in  water.  Explain  the  liberation  of 
CO2  when  HC1  is  added  to  solid  CaCO3. 

39.  Explain  why  a  knowledge  of  chemical  arithmetic  is  es- 
sential in  preparing  baking  powders. 

40.  SO2  is  a  reducing  agent  but  CO2  is  not.  Explain. 


1.  Make  a  study  of  the  cost  per  ounce  of  a  half-dozen  widely 
advertised  baking  powders.  Nearly  all  baking  powders  liberate 


the  same  amount  of  CO2  for  equal  weights  of  baking  powder. 
Is  the  wide  range  of  prices  for  such  baking  powders  warranted? 
Give  reasons  for  your  answers.  Why  do  such  wide  ranges  in 
prices  exist? 

2.  Question  your  mother,  or  other  members  of  your  family, 
and  friends  about  their  preferences  in  leavening  agents.  List 
the  data  you  obtain  and  give  your  own  reactions  to  such  pref- 

3.  Visit  a  soft-drink  bottling  plant  in  your  neighborhood, 
and  report  on  the  method  used  in  charging  the  soft  drinks 
with  CO2. 

4.  After  consulting  your  teacher  of  biology  on  sources  of 
information,  prepare  a  two-  or  three-page  report  on  micro- 
organisms, other  than  yeast,  that  are  used  by  man  in  the  prep- 
aration of  his  food  supply.  Arrange  a  class  discussion  on  this 

5.  Construct  a  glass  model  of  an  acid-sodium  bicarbonate 
type  of  portable  fire  extinguisher.  Demonstrate  its  action  be- 
fore the  class. 



.  .  .  The  introduction  of  gas  light- 
ing  (1810)  met  with  much  opposi- 
tion, as  one  can  understand,  on  the 
part  of  the  tallow  chandlers  and  sell- 
ers of  whale  oil  —  at  that  time  used 
as  an  illuminant;  but  after  a  time 
the  new  method  of  lighting  was  wel- 
comed. Alexander  Findlay,  1930 

Priestley  discovers  an  oxide  of  carbon  that  burns.  Almost  half  a 
century  after  the  discovery  of  carbon  dioxide,  another  quite  differ- 
ent oxide  of  carbon  was  found,  carbon  monoxide,  CO.  Not  until  the 
last  year  of  the  eighteenth  century  was  carbon  monoxide  known  to 
be  a  new  compound  of  carbon. 

Priestley  again  was  the  first  to  recognize  clearly  a  new  gaseous 
compound  in  the  flickering  blue  flame  which  played  over  his  fur- 
nace fire.  This  valuable  discovery  was  made  in  America,  where  this 
dissenting  English  minister  had  fled  to  escape  political  and  religious 

How  carbon  monoxide  is  formed.  Priestley's  observation  indicated 
that  when  coal  or  other  carbonaceous  fuel  is  burned  in  a  limited  sup- 
ply of  air,  carbon  monoxide,  rather  than  carbon  dioxide,  is  formed. 

2C  +  O2  ->  2CO 

Carbon  monoxide  is  formed  also  in  a  house  furnace  by  the  reduc- 
tion of  carbon  dioxide  gas  as  it  passes  over  glowing  coal,  which,  as 
you  know,  is  an  excellent  reducing  agent. 

CO2  +  C  -»  SCO 


Normally,  carbon  monoxide  burns  at  the  surface  of  a  coal  fire, 
giving  off  heat  and  changing  to  carbon  dioxide,  but  if  the  supply  of 
air  is  insufficient,  or  if  the  flame  is  chilled,  some  of  the  carbon  mon- 
oxide may  escape  without  burning. 

How  carbon  monoxide  is  prepared  in  the  laboratory.  Formic  acid 
is  a  colorless  liquid  with  a  characteristic  odor,  that  is,  a  distinct  odor 
that  can  be  used  to  identify  it.  Its  formula  is  HCOOH  (only  the 
last  H  is  a  replaceable  hydrogen  atom) .  When  warm  concentrated 
sulfuric  acid,  a  powerful  dehydrating  agent,  is  added  to  formic  acid, 
it  extracts  the  HOH  of  the  acid,  leaving  pure  CO. 

HCOOH  ->  HOH  +  CO  t 

Physical  properties  of  carbon  monoxide.  Carbon  monoxide  is  a 
colorless,  tasteless,  almost  odorless  gas,  slightly  lighter  than  air,  and 
slightly  soluble  in  water.  Its  odor  is  very  difficult  to  detect.  It  can  be 
liquefied,  but  only  under  high  pressures  and  low  temperatures. 

Chemical  properties  of  carbon  monoxide.  Carbon  monoxide  burns 
in  air  with  a  pale  blue  flame,  forming  carbon  dioxide,  which,  because 
it  is  unable  to  unite  with  any  more  oxygen,  cannot  burn. 

2CO  +  02  -»  2C02 

Because  of  carbon  monoxide's  great  attraction  for  oxygen,  it  is 
an  excellent  reducing  agent,  and  therefore  it  is  used  in  the  extraction 
of  iron  from  its  oxide.  The  carbon  monoxide  is  formed  by  the  reac- 
tion of  coke  (carbon)  and  CO2  in  the  upper  part  of  a  blast  furnace. 

Fe2O3  +  SCO  ->  2Fe  +  3CO2  T 

Carbon  monoxide  combines  with  certain  metals,  forming  a  series 
of  compounds  called  carbonyls.  For  example,  when  nickel  is  heated 
in  the  presence  of  carbon  monoxide  to  a  temperature  of  about 
40°C.,  nickel  carbonyl,  a  poisonous  liquid,  is  formed. 

Ni  +  4CO  ->  Ni(CO)4  (nickel  carbonyl) 

Upon  further  heating,  carbonyls  liberate  CO  and  give  up  their  met- 
als. Nickel  carbonyl  is  used  in  the  extraction  of  nickel  from  its  ores. 
Carbon  monoxide  in  the  presence  of  suitable  catalysts  also  com- 
bines directly  with  hydrogen,  forming  wood  alcohol  (see  page  548) , 
and  with  chlorine,  forming  phosgene. 

CO  +  2H2  -»  CH3OH  (wood  alcohol) 
CO  +  C12  — >  COC12  (phosgene) 


Some  uses  of  carbon  monoxide.  Carbon  monoxide  is  an  excellent 
fuel,  burning  with  intense  heat  and  leaving  no  solid  residue.  It  is 
most  often  used  in  conjunction  with  other  gases  as  a  mixed  fuel,  as 
in  water  gas,  for  example,  which  is  a  mixture  of  carbon  monoxide 
and  hydrogen. 

Carbon  monoxide  is  used  widely  in  the  metallurgy  (extracting  a 
substance  from  its  ores)  of  iron  and  nickel. 

Deadly  carbon  monoxide,  a  cumulative  poison.  Carbon  monoxide 
is  the  most  widespread  poison  connected  with  human  life  and  activ- 
ity. It  acts  painlessly.  Furthermore,  it  gives  no  warning  of  danger, 
because  it  is  almost  odorless,  and  its  victim  slowly  passes  into  a  state 
of  unconsciousness.  Its  physiological  action  is  caused  by  the  strong 
affinity  of  carbon  monoxide  for  hemoglobin,  an  attraction  that  is  300 
times  greater  than  the  attraction  of  oxygen  for  hemoglobin. 

Hemoglobin  is  present  in  the  red  blood  corpuscles  and  its  function 
is  to  unite  chemically  with  oxygen  and  carry  it  to  all  parts  of  the  body 
by  the  circulation  of  the  blood.  By  forming  a  stable  compound  with 
the  hemoglobin  in  blood,  carbon  monoxide  prevents  oxygen  from 
reaching  millions  of  cells  in  the  body.  The  victim  is  thus  killed  by 
suffocation  —  a  lack  of  oxygen. 

The  formation  of  carbon-monoxide-hemoglobin  takes  place  gradu- 
ally, and  this  substance  remains  in  the  blood.  When  the  percentage 
of  this  stable  compound  has  reached  40  percent,  the  victim  collapses; 
when  it  reaches  80  percent,  death  occurs.  Such  a  poison,  which  col- 
lects slowly  in  the  body  and  over  a  long  period  of  time  until  it  reaches 
a  dangerous  concentration,  is  said  to  be  a  cumulative  poison. 

The  treatment  for  carbon  monoxide  poisoning  consists  of  induc- 
ing respiration  by  the  Schafer  prone-pressure  or  the  more  recently 
adopted  Holgar-Nielsen  "back-pressure,  arm-lift"  method  in  the  open 
air,  while  carbogen  is  administered.  No  alcohol  should  be  given.  As 
you  know,  the  carbon  dioxide  makes  the  patient  breathe  more  deeply. 
The  injection  into  the  blood  stream  of  methylene  blue,  a  dye,  as  an 
antidote  for  carbon  monoxide  poisoning  has  been  reported  to  be 
effective  in  certain  cases. 

Where  the  danger  of  carbon  monoxide  poisoning  lurks.  The  ex- 
haust gas  of  automobiles,  in  which  carbon  monoxide  exists  in  vary- 
ing percentages  up  to  about  12  percent,  is  one  of  the  most  common 
sources  of  carbon  monoxide.  Hundreds  of  persons  are  killed  by  this 
gas  every  year,  usually  in  some  such  ways  as  these. 

1)  A  motorist  drives  his  car  into  his  garage.  It  is  cold,  and  he  shuts 
the  doors  of  his  garage  to  keep  warm.  He  keeps  the  motor  running 
while  he  gets  under  the  machine  to  make  some  repairs.  Before  he  is 



aware  of  it,  carbon  monoxide  from  the  exhaust  pipe  of  the  automo- 
bile causes  unconsciousness.  Death  soon  follows. 

2)  On  a  very  cold  morning  a  driver  may  start  the  engine  before 
opening  the  garage  door,  or  he  may  sit  in  a  closed  car  with  the  en- 
gine running  while  waiting  for  a  friend.  Carbon  monoxide  gas  from 
the  exhaust  finds  its  way  into  the  confined  space  and  takes  its  toll. 

The  carbon  monoxide  comes  from  the  incomplete  combustion  of 
the  gasoline.  Although  gasoline  is  a  mixture  of  hydrocarbons,  we  may 
represent  its  composition  by  the  formula  C7H16.  During  incomplete 
combustion,  the  following  reaction  is  one  that  may  take  place: 

C7H]6  +  8O2  -»  6CO  +  8H2O  +  CO2 

The  most  serious  industrial  cases  of  carbon  monoxide  poisoning 
occur  from  exposure  to  gases,  such  as  producer  and  blast  furnace 
gases  used  in  many  manufacturing  plants  (see  pages  359  and  407) . 

Most  poisonings  from  carbon  monoxide  outside  of  industrial  and 
auto  cases  are  caused  by  the  escape  of  gas  from  kitchen  ranges  and 
stoves,  gas  refrigerators  and  other  gas  appliances  because  of  leaks 
and  the  accidental  extinguishing  of  the  burning  gases.  A  furnace,  im- 
properly adjusted,  may  give  off  quantities  of  this  gas  and  become  a 
menace  to  life.  As  shown  below,  every  coal  stove  and  furnace  should 
have  an  adequate  circulation  of  air,  otherwise  carbon  monoxide  may 
be  formed  and  finally  reach  dangerous  concentrations. 

Firing  for  complete  combustion.  In  firing  a  coal-burning  furnace, 
the  object  is  to  obtain  as  much  heat  from  the  coal  as  possible.  This 
can  be  done  only  when  combustion  of  the  coal  is  complete,  or  nearly 
so,  with  a  minimum  of  intermediate  products,  including  carbon 
monoxide  and  free  carbon,  or  soot,  remaining  unburned.  Incom- 
plete combustion  is  dangerous  because  of  the  dangers  from  the  car- 
bon monoxide  produced.  It  is  also  wastefully  inefficient  and  robs 



Fig.  78.  Stove  in  operation. 
When  damper  A  is  open,  but 
dampers  B  and  C  are  not,  in- 
complete combustion  results. 
When  0  and  C  only  are  open, 
the  fire  burns  slowly. 




burning  CO 
reduction  of  CO2 

of  coal 



the  consumer  of  many  Btu  per  dollar  of  fuel.  In  addition,  the  great 
quantities  of  smoke  and  soot  produced  are  a  menace  to  health  and 
greatly  increase  expenditures  for  laundry  and  dry  cleaning. 

Complete  combustion  of  coal  depends  upon  adequate  quantities  of 
air  and  upon  proper  firing  methods,  that  is,  proper  placement  of  coal 
in  the  firebox.  In  adding  coal  to  a  burning  fire,  push  the  glowing  coals 
from  the  front  of  the  firebox  to  the  rear,  and  place  the  additional  coal 
in  the  space  from  which  the  coals  were  pushed.  In  this  way,  the  fire  is 
not  smothered  and  great  clouds  of  smoke  do  not  result.  The  glowing 
coals  at  the  rear  of  the  firebox  heat  the  added  coal,  combustible  va- 
pors are  formed  and  oxidized  relatively  completely  in  the  region  sep- 
arating the  new  coal  from  the  glowing  coals. 

If  the  new  coal  is  spread  completely  over  the  glowing  coals  the 
fire  is  partially  smothered.  Combustible  vapors  are  formed,  to  be 
sure,  but  for  a  long  period,  the  temperature  above  the  added  fuel  is 
lower  than  the  kindling  temperature  of  the  vapors,  and  very  small 
amounts  of  heat  are  produced. 

Proper  regulation  of  the  air  supply  is  a  most  important  factor  in 
securing  complete,  or  nearly  complete,  combustion.  If  the  supply  of 
air  is  insufficient,  incomplete  combustion  results,  and  carbon  mon- 
oxide and  free  carbon,  or  soot,  are  produced  in  large  quantities.  If 
too  much  air  is  available,  combustion  is  relatively  complete  but  very 
rapid,  and  results  in  uneven  heating.  If  far  too  much  air  is  available, 
much  of  the  heat  produced  is  carried  out  the  chimney  without  trans- 
ferring its  heat  to  the  water,  steam,  or  air  in  the  heating  system. 

Finding  the  damper  settings  that  result  in  relatively  complete  com- 
bustion is  not  difficult,  but  it  does  require  some  patience  and  much 
observation,  for  under  varying  outside  conditions,  damper  settings 
vary  also.  If  you  can  adjust  the  dampers  so  that  the  flame  produced 
in  the  firebox  has  a  minimum  of  yellow  flame,  the  conditions  neces- 
sary for  nearly  complete  combustion  probably  are  present. 

Coal  fires  Should  not  be  poked  or  shaken  as  often  as  most  persons 
poke  and  shake  them.  Shaking  of  ashes  should  occur  only  when  the 
accumulation  of  ashes  becomes  so  deep  that  the  air  supply  may  be  im- 
peded. Ashes  insulate  the  metal  grates  of  the  firebox  from  the  in- 
tense heat  of  the  fire,  and  if  they  are  all  removed,  the  grates  may  be 
damaged  as  a  result  of  receiving  more  heat  than  they  were  designed 
to  receive.  Poking  a  fire  frequently  spreads  the  burning  coals  through- 
out the  firebox  and  results  in  too  rapid  combustion^ 

The  simple  principles  just  discussed  will,  if  followed,  do  much 
to  increase  the  heat  that  consumers  get  from  the  c6al  they  buy;  smoke 
nuisances  will  be  reduced,  and  fuel  bills  will  be  ctil  itiaterially. 



How  carbon  monoxide  can  be  removed  from  air.  For  a  long  time 
no  simple  method  of  removing  noxious  carbon  monoxide  from  air 
was  known.  Activated  charcoal,  which  adsorbs  various  gases,  such  as 
ammonia,  acetone,  benzene,  and  chlorine,  will  not  remove  carbon 
monoxide  from  air. 

A  catalyst  called  "Hopcalite,"  which  consists  of  a  mixture  of  me- 
tallic oxides,  causes  carbon  monoxide  to  change  to  carbon  dioxide 
at  ordinary  temperatures  and  pressures.  In  the  presence  of  "Hop- 
calite," 1  molecule  of  carbon  monoxide  combines  with  another  atom 
of  oxygen  from  the  air.  A  rescue-breathing  apparatus  containing  a 
"Hopcalite"  canister  is  now  used  by  persons  who  find  it  necessary 
to  enter  regions  where  concentration  of  carbon  monoxide  is  high. 

How  can  carbon  monoxide  be  detected?  Simple  color-detector 
tubes  have  been  devised  for  determining  minute  amounts  of  carbon 
monoxide  in  air.  One  such  preparation  called  hoolamite  contains 
specially  prepared  iodine  pentoxide,  L()r>,  which  oxidizes  carbon 
monoxide  to  CO2. 

5CO  +  I2O&  -»  5CO2  +  I2 

The  iodine  fumes  liberated  cause  a  change  in  color  that  is  directly 
proportional  to  the  amount  of  carbon  monoxide  present.  This 
amount  can  be  determined  by  comparison  with  a  standard  color 
scale.  The  color  fades,  and  the  tubes  may  be  used  again.  Before  this 
method  was  developed,  canaries  were  often  used  as  detectors  because 
these  birds  are  very  sensitive  to  minute  amounts  of  this  poison  and 
show  its  effects  before  man  does. 

Official  U.S.  Navy  photograph 

Navy  firefighters  ready  for 
action.  Their  rescue-breathing 
apparatus  will  protect  them 
from  carbon  monoxide  and 
other  dangerous  fumes. 

These  recording  devices  keep 
a  constant  check  on  the  con- 
dition of  the  air  in  the  Lincoln 
Tunnel.  The  8000-foot  vehi- 
cular tunnel  connects  the 
New  York  and  New  Jersey 
banks  of  the  Hudson  River. 

The  Port  of  New  York  Authority 

In  some  tunnels,  such  as  the  Holland  and  Lincoln  Tunnels  at 
New  York  City,  machines  have  been  installed  which  record  on  a 
time  chart  the  amount  of  carbon  monoxide  present  in  the  air.  A 
warning  bell  is  caused  to  ring  when  the  carbon  monoxide  reaches 
a  concentration  of  four  parts  in  10,000.  At  this  point,  also,  dampers 
and  fans  begin  to  operate  automatically  to  change  the  air. 

What  are  the  more  important  gaseous  fuels?  The  fuel  gases 
burned  each  year  in  this  country  are  worth  more  than  two  billion 
dollars.  The  total  volume  of  gas  produced  is  about  eight  trillion 
cubic  feet.  Besides  pure  carbon  monoxide,  which  is  seldom  used 
alone,  the  most  important  gaseous  fuels  are  water  (or  synthesis)  gas, 
producer  gas,  coal  gas,  natural  gas,  acetylene  gas,  and  hydrogen.  Hy- 
drogen has  already  been  discussed;  a  brief  discussion  of  the  others 

A  gas  from  steam  and  carbon.  Water  gas  is  manufactured  by  pass- 
ing water  in  the  form  of  steam  over  glowing  coke  or  hard  coal.  The 
coke  is  first  burned  in  a  draft  of  air  until  it  is  red-hot.  The  air  is 
then  shut  off  and  the  steam  turned  on.  The  temperature  of  the  coke 
falls  gradually  because  the  reaction  is  endothermic,  and  when  it 
reaches  about  1000°C.,  the  steam  is  cut  off  and  the  air  supply  re- 
newed. During  the  process  the  carbon,  which  is  an  ideal  reducing 
agent,  combines  with  the  oxygen  of  the  water,  leaving  hydrogen. 

H2O  +  C  ->  CO  +  H2 

water  gas 

At  a  temperature  of  about  1000°C.,  the  most  gas  is  produced. 
Above  this  temperature,  the  carbon  monoxide  formed  may  react 
with  the  steam  and  change  to  carbon  dioxide. 

H2O  +  CO  ->  CO2  +  H2 



Water  gas  is  thus  a  mixture  of  carbon  monoxide  and  hydrogen. 
It  is  used,  either  alone  or  mixed  with  coal  gas,  for  domestic  heating 
purposes.  Since  both  carbon  monoxide  and  hydrogen  burn  with  al- 
most colorless  flames,  a  mixture  of  the  two  cannot  be  used  for  illumi- 
nating purposes  unless  it  is  made  luminous  by  injecting  into  it  gas- 
eous hydrocarbons  from  petroleum,  which,  on  burning,  give  a  yellow 
flame.  This  process  of  adding  hydrocarbon  vapors  to  water  gas  is 
called  enriching  the  gas.  When  water  gas  burns,  it  forms  water  va- 
por and  carbon  dioxide. 

CO  +  H2  +  O2  -» H2O  +  CO2  T 

water  gas 

When  used  in  the  synthesis  of  gasoline  and  other  chemicals  it  is 
called  synthesis  gas. 

The  gasification  of  coal.  In  areas  where  natural  gas  is  not  available 
as  a  fuel  for  factories  and  homes,  producer  gas  is  most  generally 
used.  It  is  produced  by  burning  low-grade  coal  in  a  furnace  with  a 
limited  supply  of  air.  The  chief  product  formed  during  the  incom- 
plete combustion  of  the  coal  is  carbon  monoxide.  The  gas  issuing 
from  this  furnace  is  mixed  with  large  quantities  of  nitrogen,  which 
is  too  inactive  to  unite  with  the  coal.  Producer  gas,  then,  is  chiefly 
a  mixture  of  about  60  percent  nitrogen  and  30  percent  carbon  mon- 
oxide. It  contains  also  about  ten  percent  hydrogen. 

2C  +  air  (02  +  N2)  -»  2CO  +  N2 

Producer  gas  is  a  much  poorer  fuel  than  water  gas  because  it  con- 
tains such  a  large  amount  of  nitrogen,  an  incombustible  gas.  Its 
manufacture  constitutes  the  most  efficient  means  of  converting  low- 
grade  coal  into  power. 

A  gas  similar  in  composition  to  producer  gas  was  first  proposed  by 
Mendeleyeff  and  later  produced  in  1933  in  the  Soviet  Union  by 
burning  coal  underground.  Instead  of  mining  the  coal  in  the  usual 
way,  a  coal  seam  was  sealed  off,  set  on  fire,  and  the  gases  produced 
were  brought  to  the  surface  through  pipes.  Later,  pure  oxygen  was 
mixed  with  the  blast  of  air  sent  into  the  mine,  and  the  gas  that  came 
out  of  the  mine  was  richer  in  composition  than  producer  gas. 

This  gasification  of  coal  underground  (gas  mining) ,  held  feasible 
by  the  U.S.  Bureau  of  Mines  after  several  years  of  experimentation, 
may  turn  out  to  be  a  revolutionary  development.  By  this  method  we 
may  be  able  to  recover  valuable  coal  in  mines  abandoned  because  of 
the  low-grade  nature  of  the  coal  or  because  of  the  thinness  of  the  coal 


Fig.    79.    Underground    gasiflcc 
tion  of  coal. 




vein.  In  addition,  this  method  makes  coal  working  safer  and  health- 
ier, a  consideration  which  is  always  desirable.  Another  advantage  is 
pipeline  distribution  of  this  fuel. 

Natural  gas,  the  Cinderella  fuel.  More  than  80  percent  of  the 
gaseous  fuel  consumed  each  year  in  the  United  States  is  natural  gas. 
In  various  sections  of  the  world,  especially  where  coal  and  petroleum 
deposits  are  found,  natural  gas  issues  from  the  earth  when  porous 
rocks  saturated  with  it  are  tapped.  The  gas  may  flow  out  under  pres- 
sure, or  pumping  may  be  required. 

Many  gas  wells  do  not  yield  oil,  but  an  oil  well  almost  always 
produces  both  gas  and  oil.  The  gas  from  wells  that  do  not  yield  oil  is 
usually  very  rich  in  methane,  CH4,  some  wells  yielding  as  high  as 
95  percent  methane. 

In  the  early  days  of  the  petroleum  industry,  little  or  no  use  was 
found  for  natural  gas,  and  most  of  the  wells  were  ignited  and  allowed 
to  burn  for  years.  When  the  waste  of  natural  resources  involved  in 
such  practices  was  realized,  controls  were  established  and  the  burn- 
ing of  natural  gas  greatly  diminished. 

Methane,  also  called  marsh  gas,  is  a  colorless,  practically  odorless, 
insoluble  gas  which  burns  with  an  almost  colorless  flame,  forming 
CO2  and  water  vapor. 

CH4  +  2O2  ->  CO2  +  2H2O 

It  has  a  calorific  value  about  twice  that  of  manufactured  gas  which 
it  has  largely  replaced.  It  is  about  half  as  heavy  as  air.  Natural  gas  is 
purified,  some  high  quality  casing-head  gasoline  being  obtained  in 



this  process.  H2S  is  also  removed  before  it  is  sent  into  pipelines.  It  is 
also  stripped  of  its  heavier  hydrocarbons  such  as  butane  and  propane 
before  being  piped  many  hundreds  of  miles  to  supply  factories,  farms, 
and  homes  with  light,  heat,  and  power.  For  example,  there  is  an 
1840-mile  pipeline  (the  Big  Inch)  from  Texas  to  New  York. 

Probably  the  origin  of  our  great  supply  of  natural  gas  is  the  same 
as  that  of  petroleum,  which  most  scientists  believe  to  be  the  result  of 
the  incomplete  decomposition  of  vegetable  or  animal  matter,  either 
with  or  without  bacterial  action.  The  formation  ,of  petroleum  does 
not  require  millions  of  years  as  was  formerly  believed.  This  is  demon- 
strated by  the  formation  at  the  present  time  of  petroleum  in  off-shore 
marine  sediments.  Some  scientists  believe,  however,  that  natural  gas 
originates  from  the  interaction  of  metallic  carbides  and  water,  just  as 
acetylene  is  formed  by  the  reaction  between  calcium  carbide  and 

CaC2  +  2H20  -»  Ca(OH)2  +  C^  t  (acetylene) 

A  gas  from  a  gray  solid.  The  gaseous  fuel  called  acetylene,  C2H2, 
is  colorless  and  odorless  when  pure,  very  slightly  soluble  in  water, 
and  somewhat  toxic.  Acetylene  has  a  tendency  to  explode  when  lique- 
fied. The  gas  is  therefore  not  liquefied,  but  is  forced  at  low  pressure 
into  a  solvent  called  acetone,  a  colorless  liquid  obtained  from  the 
destructive  distillation  of  wood,  and  mixed  with  some  inert  porous 
material,  such  as  wood  charcoal  or  asbestos  fiber.  It  is  sold  in  port- 
able steel  cylinders  and  is  used  widely  in  oxyacetylene  torches.  When 
the  valve  of  one  of  these  cylinders  is  opened,  the  pressure  is  reduced, 
and  some  of  the  gas  escapes  from  solution. 

For  emergency  use,  and  in  lighthouses  and  isolated  districts  where 
electric  lighting  and  illuminating  gas  are  not  available,  special  acety- 
lene generators  have  been  constructed.  These  special  generators  al- 
low water  to  come  in  contact  with  calcium  carbide  at  regulated  rates 
of  speed,  so  that  the  gas  may  be  liberated  as  needed. 





Fig.  80.  (left)  Carbide-to-water  acetylene  generator  for  large  installations,  (center) 
Water-to-carbide  generator,  (right)  The  acetylene  burner. 

The  burner  used  with  acetylene  gas  must  be  specially  constructed 
to  permit  the  access  o£  a  large  amount  of  air  to  the  burning  gas. 
Because  acetylene  is  rich  in  carbon,  it  would  otherwise  burn  incom- 
pletely and  produce  a  smoky  flame. 

2C2H2  +  5O2  —>  4CO2  +  2H2O  (complete  combustion) 
2Q>H2  +  O2  — >  4C  -h  2H2O  (incomplete  combustion) 

Alcohol,  acetone,  vinegar,  chloroform,  plastics,  and  synthetic  fibers, 
rubber,  and  gasoline  have  been  built  up  chemically  from  acetylene 
gas  (see  Chapters  34  and  35) . 

A  gas  from  the  destructive  distillation  of  coal.  About  the  time 
that  Priestley  was  studying  carbon  monoxide,  coal  gas  was  intro- 
duced as  an  illuminant.  William  Murdock,  a  Scottish  workman  em- 
ployed by  James  Watt,  developer  of  the  steam  engine,  carried  out 
experiments  that  led  to  lighting  part  of  the  Boulton  and  Watt  fac- 
tory in  Birmingham,  England,  with  gas  in  1798.  A  few  years  later, 
gaslighting  was  introduced  in  the  United  States.  Much  opposition 
was  raised  against  it,  but  the  advance  of  science  could  not  be  stopped 
for  long.  Manufacturers  of  candles  and  whale  oil,  fought  against  il- 
luminating gas,  even  as  other  manufacturers  have  since  struggled  un- 
successfully against  other  innovations. 

Coal  gas  is  obtained  from  coal  by  destructive  distillation.  Fig.  81 
shows  the  method  of  manufacture  and  purification  in  a  byproduct 
coke  oven.  Coal  is  heated  in  a  closed  oven.  The  vapors  formed  are 
first  passed  into  the  hydraulic  main,  where  some  of  the  vapors  of 

impure  gas 

Fig.  81.  Steps  in  the  production  of  coal  gas. 




coal  tar  and  ammonia  are  condensed.  The  remaining  vapors  then 
enter  the  condensers,  where  the  rest  of  the  ammonia  is  absorbed  and 
the  coal  tar  condensed.  The  purifier,  containing  iron  oxide  or  lime, 
removes  hydrogen  sulfide  and  other  sulfur  compounds.  The  purified 
coal  gas  then  enters  the  gas  holder.  Coal  gas  contains  about  ten  per- 
cent carbon  monoxide,  40  percent  hydrogen,  and  40  percent  methane. 
It  contains  also  about  ten  percent  nitrogen. 

The  destructive  distillation  of  one  ton  of  coal  yields  approxi- 
mately 10,000  cubic  feet  of  coal  gas,  20  gallons  of  ammonia  water, 
and  120  pounds  of  coal  tar.  About  1400  pounds  of  coke  remain  in 
the  retort.  A  tarry  matter  called  pitch  is  left  also,  which  is  used  as  a 
binder  in  road  construction. 

What  is  an  explosive  mixture?  All  the  gases  discussed  in  this 
chapter  burn  quietly  in  air.  However,  when  they  are  mixed  with  air 
in  the  right  proportions  to  secure  nearly  complete  combination  and 
then  ignited,  very  rapid  oxidation  takes  place,  suddenly  producing 
extremely  large  volumes  of  gases.  The  high  temperature  of  the  re- 
action helps  to  account  for  the  large  volume  of  gases  formed,  since 
gases  expand  as  their  temperature  is  raised.  High  temperatures  also 
result  in  high  pressures.  If  the  explosive  mixture  is  confined,  as,  for 
example,  in  a  mine  or  factory,  great  destruction  takes  place. 

Gases  and  vapors  differ  from  one  another  in  the  range  of  composi- 
tion of  their  explosive  mixtures.  Thus  coal  gas  will  explode  when 
anywhere  from  six  percent  to  29  percent  of  it  is  mixed  with  air.  Air 
containing  less  than  six  percent  or  more  than  29  percent  of  coal  gas 
will  not  explode. 

The  character  of  the  explosive  mixture  is  of  prime  importance  in 
the  working  of  internal-combustion  engines.  For  example,  the  degree 
of  smoothness  and  of  power  in  the  running  of  an  automobile  engine 
depends  upon  getting  the  right  mixture  of  air  and  gaseous  fuel 
admitted  to  the  cylinders  and  ignited  at  the  proper  instant  by  an 
electric  spark.  The  closer  the  proportions  of  the  mixture  to  those 
necessary  for  securing  complete  combustion,  the  greater  the  power 
produced  and  the  less  the  waste  of  fuel. 

Because  of  the  importance  attached  to  securing  this  proper  mix- 
ture, the  carburetor,  in  which  the  mixing  is  done,  is  often  called  the 
heart  of  the  internal-combustion  engine.  This  needs  adjustment  from 
time  to  time  to  insure  highest  efficiency  (see  Fig.  82) . 

Gasoline  vapor  mixed  with  air  may  form  a  very  explosive  mixture. 
Many  serious  accidents  have  resulted  from  the  careless  use  of  gaso- 
line in  dry  cleaning  at  home.  Such  cleaning  operations  should  be 
done  in  the  open  and  away  from  any  source  of  ignition.  In  using 



mixture  of 
gas  and  air 



spark  plug 

Fig.  82.  Carburetor  and  cyl- 
inder assembly.  The  gaso- 
line-air mixture  is  ignited  in 
*he  cylinder  by  the  spark 

P|ua-  The  «neray  of  tne  ex- 

plosion  is  transformed  into 
motion  by  the  piston. 

gasoline  for  dry  cleaning,  vigorous  or  continued  rubbing  of  the  fab- 
ric should  be  avoided  lest  the  friction  ignite  the  gasoline. 

Safety  measures  against  mine  explosions.  For  centuries  the  igni- 
tion of  explosive  mixtures,  especially  in  coal  mines,  has  caused  seri- 
ous loss  of  life.  In  1556,  Agricola  («-grik'6-la)  published  a  book  on 
mining,  one  section  of  which  dealt  with  "the  ailments  and  accidents 
of  miners  and  the  methods  by  which  they  can  guard  against  these." 
But  life  was  cheap  in  those  days  and  little  was  done  to  protect  min- 
ers against  explosions.  At  the  end  of  the  eighteenth  century,  there 
came  an  emphasis  on  the  rights  of  man  and  with  it  a  new  humani- 
tarianism.  An  interest  in  occupational  accidents  and  diseases  was 
aroused.  Gases  in  mines  and  mine  ventilation  were  studied. 

Davy,  who  discovered  sodium  and  potassium,  and  made  other  im- 
portant contributions  to  chemistry,  was  among  those  interested  in 
the  plight  of  the  miners  in  England.  He  devised  a  simple  safety  lamp 
to  prevent  mine  explosions.  It  is  based  upon  two  principles:  (1)  An 
explosive  mixture  does  not  undergo  chemical  change  until  its  kin- 
dling temperature  is  reached.  (2)  Metal  surfaces  spread  heat  rapidly. 

In  the  Davy  lamp,  the  flame  is  surrounded  by  a  wire  gauze,  which 
distributes  the  heat  produced  by  the  flame  over  a  wide  area,  and  thus 
prevents  the  explosive  mixture  outside  the  lamp  from  reaching  its 
kindling  temperature.  A  lighted  match  brought  over  a  wire  gauze, 
as  shown  in  the  illustration  below,  will  set  fire  to  the  gas  above  the 
gauze,  but  the  gas  below  will  not  catch  fire.  Why?  This  is  an  illustra- 
tion of  the  principle  of  the  Davy  safety  lamp. 

In  commercial  mining,  the  Davy  lamp  has  been  replaced  by  a 
battery-operated  electric  lamp. 

gas  flame 

wire  gauze 

gas  below  its 
kindling  temperature 

Fig.  83.  Demonstration  of  the 
principle  of  the  Davy  lamp. 
Why  does  the  gas  below  the 
wire  gauze  not  catch  fire? 

In  an  underground  coal 
mine,  the  air  is  tested  for 
the  presence  of  methane 
immediately  after  coal  is 
blasted  from  the  seam. 
The  flame  of  the  safety 
lamp  turns  blue  when  the 
gas  is  present. 

National  Coal  Association 


Faraday,  Michael.  The  Chemical  History  of  a  Candle. 
"Kings'  Treasuries  of  Literature"  Series,  E.  P.  Button  8c  Co., 
New  York,  1920.  This  book  consists  of  six  lectures  delivered 
by  Faraday  before  young  boys  and  girls  at  the  Royal  Institution 
of  London  in  18(>()-1861.  Lecture  I  covers  (lames  —  sources  and 
structure.  The  book  is  a  classic,  and  though  about  100  years 
old  may  still  be  read  with  pleasure  and  profit. 

Manchester,  Harland.  New  World  of  Machines,  pp.  174- 
189.  Random  House,  New  York,  1945.  "Power  for  Tomorrow" 
is  a  fine  chapter  in  this  carefully  written  book. 

Oettingen,  W.  F.  von.  Carbon  Monoxide:  Its  Hazards  and 
the  Mechanism  of  Its  Action.  Public  Health  Bull.  No.  290, 
1944.  Supt.  of  Documents,  Washington,  D.C.  35^. 

Group  A 

1.  Where  and  by  whom  was  pure  GO  first  studied? 

2.  (a)  How  is  CO  formed  in  a  furnace?  (b)  Write  an  equa- 
tion to  illustrate,   (c)  Write  the  equation  for  the  oxidation  of 

3.  (a)  How  is  CO  usually  prepared  for  laboratory  use? 
(b)  What  is  the  function  of  the  H.,SO4  used? 

4.  Compare  the  physical  properties  of  N2  and  CO. 

5.  (a)  What  are  two  chemical  properties  of  CO  in  addition 
to  its  combustibility?    (b)  Write  equations  to  illustrate  each. 

6.  Write  balanced  equations  for  the  following.  Name  the 
products,     (a)  Action    of    chlorine    and    carbon    monoxide, 
(b)  action  of  hydrogen  and  carbon  monoxide,    (c)  action  ol 
carbon  monoxide   on   heated   nickel,   and    (d)  formation   of 



,    ,    |   ,    t   T 

7.  Discuss  four  uses  of  CO. 

8.  (a)  Explain  how  CO  acts  on  the  hemoglobin  of  the 
blood  of  a  person  breathing  it.    (b)  What  first-aid  treatment 
should  be  given  a  person  who  has  been  overcome  by  CO? 
(c)  What  is  the  function  of  the  small  amount  of  CO2  ad- 

9.  Poisoning    by    CO    often    occurs    in    closed    garages, 
(a)  Why?   (b)   What  precautions  should  be  taken  to  prevent 
this  danger? 

10.  (a)  What  gases  are  generally  adsorbed  by  the  C  used  in 
gas  masks?  (b)  What  substance  is  used  in  CO  detectors? 

11.  (a)  What  are  two  important  factors  in  firing  with  coal 
for   complete    combustion?     (b)  What    is    the   most    efficient 
method  of  firing?   (c)  Why? 

12.  In  firing  with  coal,  what  are  the  results  of   (a)  inade- 
quate air  supply?  (b)  too  much  air?  (c)  far  too  much  air? 

13.  How  can  you  tell  when  the  conditions  necessary  for 
complete,  or  nearly  complete,  combustion  are  likely  to  be 

14.  (a)  What  do  large  quantities  of  black  smoke  issuing 
from  a  chimney  indicate?   (b)  Why? 

15.  (a)  Is  black  smoke  undesirable?   (b)  Why? 

16.  (a)  What  are  two  pure  gaseous  fuels?    (b)  What  are 
three  that  are  mixtures  of  gases?    (c)  Which  of    (b)  contain 
large  quantities  of  CO? 

17.  (a)  Write  an  equation  for  the  manufacture  of  water 
gas.   (b)  What  is  "synthesis  gas"? 

18.  What  volume  of  steam  is  used  in  making  1500  cu.  ft.  of 
water  gas? 

19.  What  weight  of  coal,  containing  85  percent  C,  is  used 
in  making  2000  cu.  ft.  of  water  gas?    (One ,  ounce-molecular 
weight  of  a  gaseous  substance  occupies  22.4  cu.  ft.) 

20.  Water  gas  contains  60  percent  CO  and  40  percent  H2. 
What  volume  of  air  is  necessary  for  tlie  complete  combustion 
of  200  cu.  ft.  of  this  gas? 

21.  (a)  What  is  meant  by  enriching  water  gas?    (b)  Why 
is  it  done? 


22.  Producer  gas  contains  about  60  percent  noncombustible 
N2,  yet  it  is  used  in  great  quantities  as  a  fuel  for  gas  engines, 
(a)  How  does  this  N2  affect  its  fuel  value?   (b)  What  process 
uses  producer  gas?  Why? 

23.  Explain  the  gasification  of  coal  underground. 

24.  (a)  On  the  basis  of  present  evidence,  what  do  we  think 
was  the  source  of  natural  gas?  (b)  How  does  its  present  con- 
sumption rank  with  that  of  other  gaseous  fuels?   (c)  What  is 
its  chief  constituent? 

25.  (a)  What  are   the  properties   of  methane?    (b)  What 
other  name  has  it?   (c)  Why? 

26.  Make  a  table  showing  the  source,  composition,  and  chief 
use  of  four  gaseous  fuels. 

27.  Describe  briefly  the  stages  in  the  manufacture  of  puri- 
fied coal  gas. 

28.  (a)  What  is  destructive  distillation?  (b)  Name  the  chief 
products  of  destructive  distillation  of  bituminous  coal? 

29.  (a)  How  is  C2H2  made?    (b)  What  are  its  principal 
properties?    (c)  Why  must  a  C2H2  burner  be  so  constructed 
that  it  allows  the  access  of  large  quantities  of  air? 

30.  How  many  cubic  feet  of  air  are  used  in  burning  com- 
pletely 2500  cu.  ft.  of  C2H2? 

31.  How  would  you  identify  CO2,  H2,  CO,  N2,  and  NO? 

32.  (a)  What  happens  when  an  explosive  mixture  is  ignited? 
(b)  Why  are  explosive  mixtures  dangerous?    (c)  What  is  the 
most  important  factor  to  consider  in  the  mixing  of  gases  for 
explosive  effect? 

33.  (a)  How  is   the  explosive  mixture  in   an  automobile 
regulated?  (b)  Why? 

34-  Describe  a  situation  at  home  in  which  a  dangerous  ex- 
plosive mixture  of  gases  might  be  formed. 

35.  What  might  happen  if  the  ventilating  system  in  an 
underwater  tunnel  such  as  the  Lincoln  Tunnel  in  New  York 
City  suddenly  got  out  of  order? 

Group  B 

36.  Why  does  CO  burn,  whereas  CO2  does  not  burn? 

37.  In  making  water  gas,  what  would  happen  if  the  tempera- 
ture used  were  too  high? 

38.  Why  has  the  byproduct  furnace  replaced  the  beehive 
coke  oven? 


39.  Why  is  a  blue  flame  seen  when  fresh  coal  is  added  to 
burning  fuel? 

40.  A  gas  sample  from  a  sealed  fire  area  in  a  mine  shows 
CO,  four  percent;  O2,  ten  percent;  CH4,  seven  percent;  and  N2, 
79  percent.  Is  the  fire  bla/ing,  or  is  a  methane  explosion  pos- 
sible? Explain. 

41.  Suppose  a  sample  of  gasoline  is  half  hexane,  C6H14,  and 
half  heptane,  C7H16.  How  many  cubic  feet  of  air  are  necessary 
for  the  complete  combustion  of  20  cu.  ft.  of  this  gasoline 


1.  If  you  live  near  an  oil  or  natural-gas  field,  visit  the  gas 
wells  and  write  a  report  on  the  gas  waste,  if  any.  Do  you  think 
Federal  control  of  such  waste  desirable?  Explain. 

2.  Davy  refused  to  take  out  a  patent  on  his  miner's  lamp, 
saying,  "No,  my  good  friend,  I  never  thought  of  such  a  thing; 
my  sole  object  was  to  serve  the  cause  of  humanity,  and  it  I  have 
succeeded  I  am  amply  rewarded  in  the  gratifying  reflection  of 
having  done  so."  Write  a  report  expressing  your  views  on  this 
incident.  Cite  similar  instances. 

3.  Study  the  recently  introduced  Holgar-Nielsen  "back-pres- 
sure, arm-lift"  method  and  compare  it  with  the  Schater  prone- 
pressure  method. 

4.  Using  a  test  tube,  a  one-hole  rubber  stopper,  a  short  de- 
livery tube  and  some  dry  sawdust  or  wooden  splint,  show  that 
a  combustible  gas  can  be  obtained  Irom  the  wood  similar  to 
coal  gas  from  coal.  Consult  your  laboratory  workbook. 




.  .  .  Potassium  and  sodium  are  the 
names  by  which  I  have  ventured  to 
call  the  two  new  substances  .  .  . 
.  They  agree  with  the  metals  in  opaci- 
ty, lustre,  malleability,  conducting 
powers  as  to  heat  and  electricity,  and 
in  their  qualities  of  chemical  combi- 
nation. Their  low  specific  gravity 
does  not  appear  a  sufficient  reason 
for  making  them  a  new  class. 

Sir  Humphry  Davy,  1807 

Elements  may  be  classified  as  metals  or  nonmetals.  Thus  far  we 
have  discussed  a  number  of  elements  including  oxygen,  nitrogen, 
chlorine,  bromine,  iodine,  fluorine,  sulfur,  and  phosphorus.  Each 
of  these  elements  has  more  than  one-half  the  number  of  electrons  in 
its  outer  electron  ring  that  are  necessary  to  complete  this  ring.  Hence, 
each  is  a  borrower  of  electrons,  has  a  negative  valence,  and  has  an  ox- 
ide (or  oxides)  that  is  an  anhydride  of  an  acid.  Each  of  these  ele- 
ments is  called  a  nonmetal. 

But  the  properties  of  these  elements  are  not  characteristic  of  all 
the  elements.  Certain  elements  have  less  than  half  the  number  of 
electrons  in  their  outer  electron  rings  necessary  to  complete  these 
rings.  Hence,  each  of  these  elements  is  a  lender  of  electrons,  has  a 
positive  valence,  and  possesses  an  oxide  (or  oxides)  that  is  an  an- 
hydride of  a  base.  Each  of  these  elements  is  called  a  metal.  Of  the 
100  chemical  elements,  78  are  metals. 

A  comparison  of  sulfur  and  magnesium  will  help  to  make  clear  the 
differences  between  a  metal  and  a  nonmetal. 

This  classification  of  elements  into  metals  and  nonmetals  is  very 
old,  but  it  still  continues  to  be  of  service.  We  must  remember,  how- 
ever, that  certain  elements  behave  either  as  a  metal  or  as  a  nonmetal. 



a  typical 


a  typical 

a)  6  electrons  in  outside  ring  a)  2  electrons  in  outside  ring 

b)  borrower  of  2  electrons  b)  lender  of  2  electrons 

c)  valence  is  —  2  c)  valence  is  +2 
Fig.  84.  Structure  of  the  atoms  of  a  typical  metal  and  nonmetal. 

How  metals  occur  in  nature.  Metals  occur  either  free  (uncom- 
bined) ,  or  chemically  combined  in  compounds.  Gold,  silver,  copper, 
and  platinum  are  examples  of  metals  that  occur  in  the  free  state. 
The  oxide  of  iron,  the  fluoride  of  aluminum,  the  chloride  of  sodium, 
the  bromide  of  magnesium,  the  sulfate  of  barium,  the  carbonate  of 
lead,  and  the  phosphate  of  calcium  are  examples  of  compounds  of 
metals  that  are  found  in  the  combined  state.  All  these  compounds 
are  minerals,  inorganic  substances  of  definite  composition  found  on 
the  earth.  A  mineral  such  as  mercuric  sulftde,  HgS,  or  rock  salt, 
NaCl,  from  which  an  element,  usually  a  metal,  may  be  profitably 
extracted  is  an  ore. 

How  metals  are  extracted  from  ores.  In  mining  operations,  a 
mineral  is  first  separated  mechanically  from  the  rock,  or  gangue 
(gang) ,  with  which  it  is  mixed  by  a  process  known  as  ore-dressing. 
The  particular  process  used  depends  upon  the  differences  in  prop- 
erties between  the  ore  and  the  gangue.  Low-grade  ores  must  be  con- 
centrated to  reduce  the  cost  of  extracting  the  metal.  Various  methods 
of  ore-dressing,  such  as  hand  or  gravity  separation,  leaching,  mag- 
netic, and  flotation,  are  described  in  the  discussion  of  specific  ores. 

The  process  of  extracting  a  metal  from  its  ore  is  called  metallurgy. 
In  general,  the  metallurgy  of  any  metal  depends  upon  the  nature  and 
purity  of  the  ore,  the  properties  of  the  metal  to  be  extracted,  and 
the  cost  of  the  processes.  The  four  most  widely  used  metallurgical 
processes  are:  (1)  electrolysis;  (2)  reduction  by  means  of  carbon; 

(3)  roasting,  that  is,  heating  the  ore  to  change  it  to  its  oxide  by 
oxidation   (usually  roasting  is  followed  by  reduction  with  carbon) ; 

(4)  aluminothermy,  that  is,  reduction  with  aluminum  powder. 

The  most  widely  used  metal  obtained  by  electrolysis  is  alumi- 
num. Iron  is  the  chief  metal  obtained  by  reduction,  and  copper  and 
zinc  are  the  most  common  metals  obtained  "by  roasting  followed  by 
reduction.  Chromium  and  manganese  are  made  by  aluminothermy. 
These  four  metallurgical  processes  are  discussed  in  Connection  with 
aluminum,  iron,  copper,  zinc,  and  chromium. 




Characteristic  physical  properties  of  a  metal.  Metals  have  many 
physical  properties  in  common.  A  brief  study  of  these  characteristic 
properties  helps  us  to  realize  why  the  classification  of  elements  into 
metals  and  nonmetals  has  aided  the  development  of  chemistry. 

1)  A  metal  is  lustrous.  That  is,  a  metal  has  a  definite  shine.  The 
luster  of  gold,  silver,  nickel,  and  copper  is  well  known. 

2)  A  metal  is  malleable.  That  is,  a  metal  can  be  hammered  into 
thin  sheets  such  as  gold  leaf  and  tin  foil.  The  most  malleable  metal 
is  gold,  which  has  been  beaten  into  sheets  so  thin  that  300,000  of 
them  placed  one  above  the  other  make  a  pile  only  one  inch  thick. 
Carbon,  a  nonmetal,  breaks  when  hammered;  it  is  brittle. 

3)  A  metal  is  ductile.  That  is,  a  metal  can  be  drawn  into  wire. 
Platinum,  for  example,  can  be  drawn  into  such  a  fine  wire  that  it 
cannot  be  seen  by  the  unaided  eye.  The  degree  of  ductility  of  metals 
varies  greatly.  Sulfur,  a  nonmetal,  cannot  be  drawn  into  wire. 

4)  A  metal  is  a  good  conductor  of  heat.  Place  one  end  of  a  copper 
wire  a  foot  long  in  a  flame  and  notice  how  quickly  your  fingers  hold- 
ing the  other  end  of  the  wire  become  warm. 

5)  A  metal  is  a  good  conductor  of  electricity.  Although  silver  is 
the  best  conductor  of  electricity,  copper  is  used  most  generally  be- 
cause of  its  much  lower  cost.  On  the  basis  of  evidence  now  available, 
scientists  think  that  good  conductors  of  electricity,  including  metals 
in  general  and  copper  and  silver  in  particular,  are  composed  of 
atoms  whose  outer  rings  have  only  a  few  electrons  loosely  held.  These 

(left)  A  prospector  in  Utah  examines  a  sample  of  uranium  ore.  (right)  Testing  the 
ore  with  acid  for  the  presence  of  copper  and  limestone. 

I'hotogmph*,  Standard  Oil  Company  (Ar./.) 

Pouring  molten  copper  from  an  electric 
furnace  into  a  mold.  The  "melt"  con- 
sists of  both  virgin  metal  and  high- 
grade  scrap. 

/>'<  xeurch  Association 

electrons  can,  therefore,  move  along  to  the  next  atom,  and  so  on, 
producing  a  flow  of  electricity.  Nonmetals,  such  as  sulfur,  have  in 
their  outer  rings  a  large  number  of  electrons,  not  so  free  to  move. 
Hence  nonmetals  are  poor  conductors  of  electricity. 

6)  Other  properties  of  a  metal.  All  metals,  with  the  exception  of 
mercury,  are  solids  at  ordinary  temperatures.  They  range  in  melting 
points  from  mercury  (  —  39° C.)  to  wolfram  (3380°C.)  .  Metals  dif- 
fer widely  in  tensile  strength,  that  is,  the  ease  with  which  they  can 
be  pulled  in  two.  Some  steel  has  a  tensile  strength  of  500,000  Ibs. 
per  sq.  in.  cross  section. 

They  range  in  density  from  lithium  (a  little  more  than  half  as 
heavy  as  water)  ,  whose  density  is  0.53,  to  osmium,  whose  density  is 
22.5.  Most  metals  are  gray  in  color,  the  two  most  common  excep- 
tions being  copper  (red)  and  gold  (yellow) .  They  range  in  hard- 
ness from  lithium,  which  is  as  soft  as  wax,  to  others  which  are  very 
hard.  All  metals  are  crystalline. 

Carbide  and   Carbon   Chemical*   Company 

A  flotation  cell  in  which  the  ores 
of  non-ferrous  metals  are  con- 
centrated. Unwanted  impurities 
settle  to  the  bottom  of  the  cell 
and  the  ore  is  skimmed  off  with 
the  froth. 




Characteristic  chemical  properties  of  a  metal.  Metals  also  possess 
many  chemical  properties  in  common. 

1)  Certain  metals  unite  with  oxygen,  forming  oxides  that  are  basic 
anhydrides.  The  burning  of  magnesium,  for  example,  produces  mag- 
nesium oxide,  which  is  the  anhydride  of  magnesium  hydroxide. 

2Mg  -f  O>  -»  2M gO;     MgO  +  H2O  ->  Mg(OH)2 

2)  Certain  metals  unite  with  water,  forming  either  bases  or  oxides, 
and  liberating  hydrogen.  Thus  steam  passed  over  hot  iron  forms  iron 
oxide  arid  hydrogen,  whereas  calcium  reacts  with  water,  forming 
calcium  hydroxide  and  hydrogen. 

3Fe  +  4H2O 
Ca  +  2H2O 

>  Fe304  4-  4H2  T 

>  Ca(OH)2  +  H2 

3)    Certain  metals  decompose  acids,  liberating  hydrogen  or  other 


Zn  +  2HC1 

Cu  +  2H2SO4 

3Cu  +  8HNO3 

CuS04  +  2H20  -f  SO,  t 

>  3Cu(NOs)2  4-  4H20  -f  2NO 

4)  Metals  combine  with  nonmetals,  forming  salts.  For  example, 
they  unite  with  chlorine,  sulfur,  and  bromine,  forming  chlorides, 
sulfidcs,  and  bromides,  respectively. 

5)  Certain  metals  unite,  with  bases,  liberating  hydrogen.  Thus, 
both  aluminum  and  zinc  react  with  sodium  hydroxide,  liberating 

2A1  +  6NaOH  -*  2Na,<AlO3  (sodium  aluminate)  -f  3H2  T 
Zn  -f  2NaOH  ->  Na2  ZnO2  (sodium  zincate)  +  H2  \ 

Copper  and  Brass  Research  Association 

These  plates  of  impure 
copper  are  about  to  be  sus- 
pended in  a  tank  contain- 
ing copper  sulfate  and 
sulfuric  acid.  The  copper 
will  be  refined  by  an  elec- 
trolytic process  described 
in  Chapter  28. 


Metals  may  be  listed  in  the  order  of  their  replacement  power. 

It  is  generally  known  that  gold  does  not  tarnish  in  air  and  is  not 
acted  upon  by  any  one  acid,  It  is  equally  well  known  that  iron, 
on  exposure  to  air,  is  oxidized  readily.  We  could,  if  it  served  any 
useful  purpose,  arrange  all  the  metals  in  the  order  of  their  ability 
to  resist  oxidation. 

A  more  useful  arrangement  is  based  upon  the  ability  of  one 
metal  to  replace  another  from  a  solution  of  a  salt  of  the  latter.  For 
example,  when  an  iron  nail  is  placed  in  a  solution  of  copper  sulfate, 
the  iron  becomes  coated  with  a  layer  of  pure  copper,  and  iron  sul- 
fate is  formed. 

Fe  +  CuSO4  -»  Cu  +  FeSO4 

Similarly,  copper  placed  in  a  solution  of  silver  nitrate  becomes 
coated  with  pure  silver  and  will,  itself,  go  into  solution  as  copper 
nitrate.  The  reverse  reactions  will  not  take  place  under  normal  con- 
ditions. That  is,  a  copper  nail  placed  in  an  iron  sulfate  solution 
will  not  deposit  iron. 

If  we  try  these  experiments,  using  a  number  of  different  metals 
and  their  salts,  we  can  arrange  the  common  metals  in  a  definite 
replacement  series.  This  replacement  series  of  the  common  metals  is 
also  called  the  electrochemical  series,  and  the  electromotive  series. 

How  the  electron  theory  explains  the  replacement  series.  Metals 
differ  in  their  tendency  to  lose  electrons  and  become  ions.  When 
free  iron  (Fe°)  replaces  copper  from  a  copper  sulfate  solution,  the 
following  electron  reaction  occurs: 

loses  2  electrons 

Fe°  +  Cu++(SO4)~  -+  Fe++(SO4)--  +  Cu° 

s, ^ 

gains  2  electrons 

Electrically  neutral  iron  loses  two  electrons  and  becomes  ionic 
Fe++,  which  is  positively  charged.  Ionic  copper  from  the  copper  sulfate 
solution  gains  two  electrons,  becomes  electrically  neutral,  and  precipi- 
tates out  as  free  copper  atoms.  The  sulfate  ions  remain  unchanged. 



Zinc  ^ 



OF  THE  |  lead 

COMMON  METALS  1  ^I0*"1 










Fig.  85.  A  strip  of  metal,  placed  in  a  solution  of  a  salt  of  a  metal  below  it  in 
the  replacement  series,  replaces  the  less  active  metal,  which  precipitates  onto  the 

The  only  change  that  takes  place,  then,  is  a  transfer  of  two  electrons 
from  free  iron  to  ionic  copper.  Thus  it  seems  that  iron  has  a  greater 
tendency  than  copper  to  lose  electrons.  That  is,  iron  is  more  metallic 
than  copper,  and  hence  appears  higher  up  in  the  replacement  series. 
If  we  place  iron  in  a  solution  of  calcium  chloride,  no  reaction  takes 
place  because  the  tendency  of  iron  to  lose  electrons  is  less  than  that 
of  calcium. 

Fe°  +  Ca++  +  2C1~  -*  no  reaction 

The  elements  high  up  in  the  replacement  series  are  so  typically 
metallic  and  have  such  a  great  tendency  to  lose  electrons,  that  even 
light  causes  them  to  throw  off  electrons.  This  fact  is  made  use  of  in 
the  photoelectric  cell.  Such  a  cell  is  frequently  lined  with  a  thin  film 
of  potassium,  rubidium,  or  cesium,  the  most  active  metal  known. 
When  light  strikes  this  film,  it  throws  off  electrons,  which  travel  to 
a  positively  charged  plate  in  the  center  of  the  cell.  A  very  feeble 
electric  current  is  produced.  This  feeble  current,  whose  strength 
depends  upon  the  intensity  of  the  light  that  strikes  the  cell,  may  be 
amplified  and  thus  made  to  control  larger  supplies  of  energy.  This 
amplified  current  may  close  a  switch,  called  a  relay,  which  will  start 
a  motor  and  open  a  door,  or  count  people  going  through  a  passage. 

The  photoelectric  cell  is  a  vital  part  of  sound-motion-picture  and 
television  equipment,  and  it  is  used  also  in  the  transmission  of  pic- 
tures by  wire.  Photoelectric  cells  are  used  in  controlling  the  "blow" 
of  Bessemer  converters  in  making  steel. 

Selenium,  an  element  belonging  to  the  sulfur  family,  was  used  in 
certain  "electric  eyes."  Selenium  is  a  good  electric  insulator  in  the 
dark,  but  in  light  it  conducts  an  electric  current  to  some  extent. 
Later  it  was  replaced  by  copper  covered  by  a  thin  film  of  Cu,O. 

vacuated  bulb 

cesium  film 

Fig.    86.    Construction 
of  a  photoelectric  cell. 


The  replacement  series  of  the  nonmetals.  In  studying  the  halo- 
gens, we  learned  that  they,  too,  could  be  arranged  according  to 
their  ability  to  replace  one  another  from  solutions  of  their  salts. 
Thus,  when  chlorine  is  added  to  a  solution  of  sodium  bromide, 
bromine  is  liberated  and  sodium  chloride  is  formed.  According  to 
the  electron  theory,  this  reaction  is  explained  as  follows: 

gains  1  electron 

SNa+Cl"  +  Br2° 

loses  1  electron 

Chlorine  is  more  typically  a  nonmetal  than  is  bromine.  There- 
fore, it  has  a  greater  tendency  to  gain  electrons.  Free  chlorine,  which 
is  neutral,  takes  one  electron  from  the  bromine  ion  and  changes  to 
Cl~,  that  is,  it  goes  into  solution.  The  bromine  ion,  on  the  other 
hand,  after  losing  one  electron,  becomes  electrically  neutral,  changes 
to  the  atomic  form,  joins  with  another  atom  of  bromine,  and  is  liber- 
ated as  a  free  bromine  molecule.  Other  nonmetals  also  may  be 
grouped  in  a  series  according  to  their  replacement  powers. 

The  replacement  series  of  the  metals  is  a  useful  tool.  Under- 
standing the  replacement  series  of  the  metals  is  of  great  value  in 
studying  these  elements.  Those  above  iron  in  the  series,  are  so  very 
active  that  they  are  never  found  free,  while  those  below  iron  occur 
both  in  the  free  state  and  chemically  combined.  As  we  go  down  the 
series,  the  tendency  of  the  metal  to  lose  electrons  diminishes,  and 
hence  the  tendency  to  oxidize  and  to  react  with  water  or  acids 
diminishes  also.  Thus  gold  and  platinum,  which  are  at  the  bottom 
of  the  list,  do  not  oxidize  in  air  even  when  hot,  and  are  not  attacked 
by  water  or  even  by  any  one  acid.  Therefore,  knowing  the  position 
of  a  metal  in  this  series,  you  can  predict  fairly  well  its  chemical  prop- 

Why  hydrogen  is  included  in  the  replacement  series  of  the  metals. 
Hydrogen,  which  from  its  physical  properties  could  never  be  con- 
sidered a  metal,  belongs  in  the  replacement  series  because  its  ion  is 
usually  positively  charged  and  behaves  chemically  as  a  metal.  All 
metals  above  hydrogen  in  the  replacement  series  liberate  hydrogen 
from  acids.  Those  metals  below  hydrogen  require  oxidizing  acids 
to  dissolve  them  and  liberate  gases  other  than  hydrogen  from  these 
oxidizing  acids.  For  example,  when  copper  reacts  with  nitric  acid, 
nitrogen  dioxide  is  given  off. 

Humphry  Davy  isolates  potassium  and  sodium.  At  the  head  of 
the  replacement  list  of  the  metals  are  the  alkali  metals,  lithium, 

Sir  Humphry  Davy  (1778-1829),  the  son  of 
a  poor  wood  carver,  was  born  at  Penzance, 
Cornwall.  The  poet,  Samuel  Coleridge,  de- 
clared that  if  Davy  "had  not  been  the  first 
chemist,  he  would  have  been  the  first  poet 
of  his  age." 

National  Portrait  Gallery,  London 

potassium,  and  sodium.  Two  of  these  elements  had  been  known  for 
a  long  time  as  part  of  the  alkaline  compounds,  potash  and  soda,  be- 
fore they  were  isolated  in  a  pure  state.  The  alkali  metals  were  orig- 
inally found  in  the  ashes  of  certain  plants.  The  name  was  taken  from 
the  Arabic  al,  meaning  the,  and  quili,  meaning  ashes. 

Because  of  the  extreme  activity  of  these  metals,  many  unsuccessful 
attempts  had  been  made  to  isolate  them.  After  the  discovery,  in  1800, 
of  the  galvanic  current  and  the  invention  of  the  electric  battery 
which  soon  followed,  scientists  used  this  new  force  in  an  effort  to 
isolate  sodium  and  potassium. 

Humphry  Davy,  an  eminent  English  chemist,  rose  from  humble 
beginnings  to  knighthood.  He  was  the  first  to  prove  chlorine  an 
element  and,  incidentally,  he  was  versatile  enough  to  spend  his  lei- 
sure hours  writing  fairly  good  poetry.  In  1807  he  sent  the  energy  of 
150  electric  cells  through  molten  potassium  hydroxide.  At  the  nega- 
tive platinum  electrode,  Davy  saw  globules  of  a  silvery  substance 
form,  and  then  spontaneously  catch  (ire.  "His  joy  knew  no  bounds, 
he  began  to  dance,  and  it  was  some  time  before  he  could  control 
himself  to  continue  his  experiments." 

London  received  Davy's  isolation  of  potassium  as  another  wonder 
o£  the  world,  and  he  was  lionized.  Some  people  paid  100  dollars  to 
attend  his  lectures  on  chemistry.  Soon  afterward,  Davy  obtained  free 
sodium  in  the  same  way,  and  lithium  from  fused  LiCl. 

How  sodium  is  prepared  for  industrial  use.  The  most  recent 
method  of  obtaining  sodium  in  large  amounts  differs  somewhat  from 
the  method  originally  used  by  Davy.  An  electric  current  is  sent 
through  melted  sodium  chloride  in  a  cell,  such  as  is  shown  in  Fig. 


sodium  collect! 


to  DC  source 


melted  NaCI 
metal  screen 
MM    head 

Fig.  87.  Downs  cell  for 
the  preparation  of  sodium. 
The  cell  was  invented  in 
1924  by  J.  C.  Downs,  an 
American  chemist. 

iron  or  copper  cathode 
«>$*    graphite  anode 

87.  Sodium  ions  travel  to  the  iron  or  copper  cathode,  gain  electrons, 
become  sodium  atoms,  and  collect  as  a  mass  of  metallic  sodium, 
which  is  drawn  off  from  time  to  time.  Chlorine  ions  travel  to  the 
graphite  anode,  lose  electrons,  and  become  gaseous  chlorine,  which 
leaves  the  apparatus  as  shown.  This  entire  process  is  continuous.  Po- 
tassium and  lithium  may  be  prepared  in  this  same  way  using  melted 
potassium  chloride  and  lithium  chloride. 

Physical  properties  of  potassium  and  sodium.  Sodium  is  a  soft, 
silvery  metal  that  melts  just  below  the  boiling  point  of  water.  Po- 
tassium, which  is  also  soft  and  silvery  white,  melts  at  an  even  lower 
temperature.  They  are  both  lighter  than  water.  Strangely  enough, 
these  two  solids  when  mixed  form  a  liquid  alloy  at  ordinary  tem- 

Some  chemical  properties  of  sodium  and  potassium.  The  electron 
pictures  of  lithium,  sodium,  and  potassium  are  shown  in  Fig.  88. 

From  these  pictures  we  can  tell  that  the  valence  of  each  of  these 
elements  is  one.  Each  reacts  with  nonmetals,  forming  salts,  and,  on 
exposure  to  the  oxygen  of  the  air,  each  is  quickly  tarnished  with  a 
coating  of  its  oxide. 

Since  sodium  has  only  one  electron  to  lend,  and  oxygen  must 
borrow  two  electrons  to  complete  its  outer  ring,  two  sodium  atoms 
combine  with  one  atom  of  oxygen,  and  the  oxide  of  sodium  is, 
therefore,  Na2O.  Sodium  peroxide,  Na.X).,,  (Na  —  O  —  O  —  Na)  ,  is 
formed  when  sodium  is  heated  in  air  free  from  carbon  dioxide. 
Because  of  its  extreme  activity,  sodium  cannot  be  kept  exposed  to 
air  or  under  water.  It  is  usually  stored  under  kerosene,  because  kero- 
sene contains  no  oxygen.  At  high  temperatures  both  sodium  and 
lithium  combine  with  hydrogen  to  form  hydrides  which  react  with 
water  liberating  hydrogen. 

LiH  +  H2O  -»  LiOH  +  H2  1 

Fig.  88.  Structure  of  the  atoms  of  the  alkali  metals. 

Lithium  .'""'"^* 

Sodium  x  C  — 

Potassium  x^- "~ 




What  has  been  said  about  the  chemical  activity  of  sodium  applies 
also  to  other  alkali  metals,  namely,  potassium,  lithium,  cesium,  and 

How  the  alkali  metals  are  used.  In  a  recent  year  150,000  tons  of 
sodium  metal  were  used  in  the  manufacture  of  several  compounds, 
such  as  sodium  peroxide,  sodium  cyanide,  sodamide  (NaNH,,)  used 
in  making  indigo,  sodium  hydride  (NaH)  used  as  a  reducing  agent 
in  removing  surface  oxides  from  steel,  and  several  detergents.  A 
sodium-lead  alloy  is  used  in  the  manufacture  of  tetraethyl  lead. 
Liquid  sodium,  because  it  is  an  excellent  heat  conductor,  is  em- 
ployed as  a  coolant  in  some  nuclear  reactors.  Sodium  is  also  used  in 
the  hot  cathode  sodium  vapor  lamp,  which  gives  twice  as  much 
light  as  the  common  filament  electric  lamp  using  the  same  amount 
of  current.  This  type  of  lamp  is  used  chiefly  in  outdoor  lighting. 

How  can  we  test  for  the  ions  of  sodium,  potassium,  and  lithium? 
If  a  clean  platinum  wire  is  dipped  into  a  salt  of  potassium  and  then 
placed  in  a  nonluminous  bunsen  (lame,  the  flame  becomes  violet  in 
color.  The  flame  of  all  sodium  salts  is  a  distinct  yellow;  that  of  all 
lithium  salts  is  red;  cesium  gives  a  bright  blue  flame.  Since  the  pres- 
ence of  even  a  trace  of  a  sodium  salt  will  obscure  the  violet  color 
of  potassium,  the  flame  of  a  potassium  salt  can  frequently  be  de- 
tected only  when  viewed  through  a  piece  of  blue  cobalt  glass,  which 
absorbs  yellow  light  rays. 

(left)  Removing  a  sodium  brick  from  a  shipping  drum,  (right)  Making  sodium 
pellets  for  laboratory  use.  Note  the  protective  equipment  in  each  case. 

Ethvl  Corporation 

The  bright  line  spectrum  of  lubricating  oil  (left)  and  of  low-grade  coal. 

Sucli  flame  tests  are  used  on  other  metals  besides  the  alkali  metals. 
Thus,  heated  copper  imparts  a  green  color  to  the  flame,  and  calcium 
gives  the  flame  an  orange-red  coloration. 

How  is  a  spectroscope  used?  In  1854,  David  Alter,  a  Pennsylvania 
physician,  described  a  method  of  detecting  an  element  by  the  color 
that  it  imparts  to  a  flame.  He  also  predicted  the  use  of  this  method 
in  determining  the  presence  of  elements  in  the  sun. 

Five  years  later,  Bunsen  and  Kirchhoff  (kirK/hof)  devised  an  in- 
strument called  the  spectroscope,  which  has  since  become  a  very 
powerful  tool  in  the  hands  of  chemists,  physicists,  and  astronomers. 
In  I860,  with  the  aid  of  this  instrument,  two  new  elements,  cesium 
and  rubidium,  were  detected  by  Bunsen  in  a  few  grams  of  salt  ob- 
tained by  the  evaporation  of  40  tons  of  spring  water. 

Nine  years  later,  the  element  helium  was  discovered  with  the  aid 
of  the  spectroscope  by  Janssen  and  Lockyer  independently.  Helium 
was  found  not  on  the  earth,  but  in  the  sun,  more  than  90  million 
miles  away.  Before  the  close  of  the  century,  this  new  element  was 
found  on  our  own  planet  by  Ramsay. 

With  the  aid  of  the  spectroscope,  other  elements,  present  in  such 
minute  quantities  that  they  had  heretofore  escaped  discovery  by 
even  the  most  delicate  instruments  of  science,  were  finally  brought 
to  light.  Today,  the  spectroscope  is  used  also  in  the  study  of  the 
complex  structure  of  the  atom. 

This  spectfOf  rctpft  makes  Him  records  of  spec- 
tra, simitar  to  shown  «bove»  From  tuch 
tfMKtra,  the  composition  of  compounds  may 

An  analytical  chemist  using  elec- 
tron diffraction  equipment  to 
identify  crystalline  material  such 
as  nickel  oxide.  Beams  of  elec- 
trons aimed  at  the  sample  break 
into  a  pattern  which  is  made 
visible  by  the  equipment. 

The  principle  of  the  spectroscope.  To  understand  the  principle 
of  the  spectroscope,  we  must  understand  a  few  things  about  light. 
Isaac  Newton,  in  1672,  performed  a  classic  experiment.  He  let  a 
beam  of  sunlight  pass  through  a  narrow  slit  into  a  dark  room  and 
placed  a  glass  prism  in  its  path.  A  band  of  colors  called  the  spectrum 
was  formed.  This  can  be  explained  by  remembering  that  light  is  a 
form  of  energy  which  is  transmitted  by  waves. 

Sunlight  is  made  up  of  light  of  various  colors.  Each  color  has  a 
different  wave-length.  Red  has  the  longest  wave-length  (0.0000(i8 
cm.) ,  and  violet  has  the  shortest  wave-length  (0.000040  cm.) .  The 
glass  prism  bent  and  split  up,  or  refracted,  the  sunlight.  The  light 
which  was  refracted,  or  bent,  least  was  the  red,  and  that  which  was 
refracted  most  was  the  violet.  Study  the  spectrum  shown  in  the  illus- 
tration in  color  following  page  382. 

Though  some  self-luminous  sources,  such  as  the  sun,  have  a  con- 
tinuous spectrum,  as  shown  in  the  illustration,  an  incandescent  va- 
por or  gas,  such  as  heated  sodium  vapor  or  electrified  neon  gas,  has 
a  discontinuous,  or  bright-line)  spectrum.  The  glowing  vapor  of  each 
element  has  its  own  characteristic  colored  band  of  light.  Thus,  so- 
dium vapor  has  one  bright  yellow  line,  lithium  has  one  red  and  one 
yellow  line,  and  the  vapor  of  iron  has  several  hundred  lines. 

The  spectroscope,  the  most  essential  part  of  which  is  a  glass  prism, 
makes  possible  the  quick  analysis  of  incandescent  vapors,  and  the 
detection  of  the  smallest  trace  of  an  element.  Less  than  a  millionth 



of  a  milligram  of  sodium,  and  minute  traces  of  poisons  in  blood 
can  be  detected  by  spectroscopic  analysis.  A  spectroscope  is  thus  a 
most  useful  tool  to  the  chemists  who  specialize  in  the  analysis  of 
many  kinds  of  substances.  Such  specialists  are  called  analytical  chem- 

Spectroscopic  work  of  great  precision  is  carried  on  today  by  means 
of  the  spectrograph.  This  instrument  differs  from  an  ordinary  spec- 
troscope in  that  the  observing  telescope  is  replaced  by  a  camera, 
which  makes  a  photographic  record  of  the  spectrum  under  examina- 
tion. This  permanent  photographic  record  makes  possible  a  more 
careful  analysis  of  the  spectrum.  Spectographic  analysis  is  used  in 
steel  and  other  alloys  and  compares  favorably  with  the  usual  routine 
quantitative  analysis.  Analysis  of  matter  by  the  study  of  spectra  is 
called  spectrum  analysis.  Because  chemical  manipulations  are  un- 
necessary, all  measurements  may  be  made  quickly. 


Chemistry,  Jan.,  1945,  pp.  37-43,  and  May,  1952.  Published 
by  Science  Service,  Washington,  D.C.  Contain  the  original 
papers  describing  the  discovery  of  lithium,  sodium,  potassium, 
cesium,  and  rubidium,  and  more  information  on  the  other 

Mills,  John.  Electronics:  Today  and  Tomorrow,  pp.  84-94. 
D.  Van  Nostrand  Co.,  New  York,  1944.  Tells  the  story  of  the 
photoelectric  cell. 

Pough,  Frederick  H.  A  Field  Guide  to  Rocks  and  Minerals. 
Houghton  Mifflin  Co.,  Boston,  1953.  An  excellent  book  for  the 
boy  or  girl  interested  in  mineral-collecting  as  a  hobby.  Attrac- 
tively illustrated. 


1.  A  metal  is  an  element  that  lends  electrons.  It  has  a  posi- 
tive valence,  and  its  oxide  is  a  basic  anhydride. 

2.  An  ore  is  a  mineral  from  which  an  element  may  be 
profitably  extracted. 

3.  Ore-dressing  is  the  separation  of  a  mineral  from  the 
valueless  rock,  or  gangue,  with  which  it  is  mixed. 

4.  Metallurgy  is  the  process  of  extracting  a  metal  from  its 
ore.  The  four  chief  metallurgical  processes  are  (1)  electrolysis, 
(2)  reduction  by  means  of  carbon,    (3)    roasting,  or  heating, 


the  ore  to  change  it  to  its  oxide,  and    (4)  aluminothermy, 
using  aluminum  powder. 

5.  Metals  have  certain  characteristic  physical  properties. 
A  metal  is  (1)  lustrous,  (2)  malleable,  (3)  ductile,  (4)  a  good 
conductor  of  heat,   (5)  a  good  conductor  of  electricity,   (6)  a 
solid  —  mercury  is  an  exception,  (7)  crystalline,  and  (8)  vari- 
able tensile  strength. 

6.  Metals  have  certain  characteristic  chemical  properties. 
Some   (1)  unite  with  oyxgen,  forming  oxides  which  are  basic 
anhydrides,    (2)  unite  with  water,   forming  either  bases   or 
oxides  with  the  liberation  of  hydrogen,   (3)  decompose  acids, 
liberating  either  hydrogen  or  some  other  gas,    (4)  combine 
with  nonmetals,  forming  salts,  and   (5)  unite  with  bases,  lib- 
erating hydrogen. 

7.  Metals  may  be  so  arranged  in  a  replacement  series  that 
each  metal  in  the  list  will  replace  each  metal  below  it  from  a 
solution  of  its  salt.  According  to  the  electron  theory,  this  be- 
havior is  explained  by  the  fact  that  the  metals  at  the  top  of 
the  list  lose  electrons  more  easily  than  those  at  the  bottom, 
and  hence  go  into  solution  more  readily.  The  metals  lower 
on  the  list  take  these  lost  electrons,  become  electrically  neutral, 
and  precipitate  out  as  free  metals. 

8.  Some  nonmetals,  such  as  the  halogens,  may  also  be  ar- 
ranged in  a  replacement  series.  According  to  the  electron  the- 
ory, chlorine  has  a  greater  tendency  to  borrow  electrons  than 
the  other  halogens  and,  hence,  borrows  electrons  from  those 
below  it  in  the  series. 

9.  The  flame  test  for  identifying  a  metal  consists  of  heating 
a  metal  or  one  of  its  salts  in  a  flame  and  noticing  the  color  that 
it  imparts  to  the  flame 

10.  The  spectroscope  is  an  instrument  devised  by  Bunsen 
and  Kirchhoff.  It  is  used  to  detect  minute  traces  of  elements 
in  incandescent  vapors.  The  most  essential  part  of  a  spectro- 
scope is  a  glass  prism,  which  disperses  or  breaks  up  the  light 
into  colored  lines  which  are  characteristic  for  each  element. 

11.  Spectrum  analysis  has  been  used  in  detecting  the  pres- 
ence of  rare  elements,  such  as  the  inert  gases  of  the  atmosphere. 
It  is  also  used  in  the  study  of  the  structure  of  the  atom. 


Group  A 

1.  According  to  the  electron  theory,  how  do  metals  differ 
from  nonmetals? 


2.  What   are   four    typical   nonmetals   and   four   typical 

3.  (a)  Name  a  few  metals  that  occur  free,    (b)  Compare 
their  chemical  activity  with  the  chemical  activity  of  metals 
found  combined. 

4.  (a)  In  what  way  does  a  mineral  differ  from  an  ore? 
(b)  Give  an  example  of  each,    (c)  Do  minerals  ever  become 
ores?  (d)  Illustrate. 

5.  (a)  What  are  the  four  most  general  methods  used  in 
metallurgy?   (b)  Give  an  example  of  an  element  extracted  by 
each  of  these  methods. 

6.  What  are  the  characteristic  physical  properties  of  metals? 

7.  (a)  Arrange  these  eight  metals  in  a  replacement  series: 
Zn,  H,  Cu,  Na,  K,  Pt,  Au,  and  Pb.    (b)  On  your  list,  check 
those  that  will  replace  the  hydrogen  of  dilute  HC1. 

8.  Arrange  the  number  of  each  metal  opposite  the  letter 
of  the  property  of  which  it  is  an  outstanding  example. 

a)  Best  conductor  of  electricity  1)  Cu 

b)  Extremely  ductile  2)  Au 

c)  Most  malleable  of  all  3)  Li 

d)  Lightest  metal  4)  Hg 

e)  Heaviest  metal  5)  W 

f)  Liquid  metal  6)  Os 

g)  Extremely  high  melting  point  7)  Ag 
h)  Reddish  luster  8)  Pt 

9)    Na 

9.  (a)  Write  five  equations  illustrating  five  chemical  prop- 
erties of  metals,  (b)  In  each  case  state  the  property. 

10.  According  to  the  electron  theory,  why  does  an  iron  nail 
become  coated  with  Cu  when  placed  in  a  solution  of  CuSO4? 

11.  In  which  of  the  following  would  a  replacement  reaction 
take  place?  Complete  the  equations  for  such  replacements. 

a)  Zn  +  Hg(NO3)2-*  e)        Cu  +  AgNO3  -» 

b)  Zn  +  CuSO4->  f)          Cu  +  ZnSO4-> 

c)  Zn  +  AgNO3  -»  g)  Cu  +  Hg(NO3)2  -> 

d)  Zn  -f  Pb(NO3)2  -»    ,        h)    Cu  +  Pb(NO3)a -» 

12.  Arrange  a  replacement  sefies  of  some  nonmetals  you 
have  studied. 

13.  According  to  the  electron  theory,  explain  why  Br2  lib- 
erates free  I9  from  a  solution  of  KI. 


14.  An  element  X  is  not  found  free.  It  attacks  warm  water, 
liberating  H2,  and  tarnishes  readily  in  air.  Where  would  you 
place  it  in  the  replacement  list  of  metals?  See  Tat>le  12. 

15.  (a)  What  are  two  elements  other  than  Na  that  belong 
to  the  sodium  family?   (b)  Why  are  they  called  alkali  metals? 

16.  (a)  How  and  by  whom  were  Na  and  K  first  isolated? 
(b)  How  was  the  news  of  this  achievement  received? 

17.  (a)  What  are  three  ways  in  which  Na  is  similar  to  K? 

(b)  one  way  in  which  it  is  different? 

18.  By  a  diagram,  show  the  present  method  of  obtaining 
free  Na. 

19.  Make  a  diagram  of  the  atom  of  Li,  and,  from   this 
diagram,    state    its     (a)  valence,     (b)  chemical    activity,    and 

(c)  atomic  number. 

20.  Using  an  equation,  describe  the  action  of  Na  on  H2O. 

21.  What  weight  of  NaOH  must  be  decomposed  to  produce 
69  g.  of  pure  Na? 

22.  A  piece  of  Na  is  placed  in  H2O,  and  336  ml.  of  H2  are 
collected.   What   weight   of   Na   took   part   in   this   chemical 

23.  Which  will  require  more  H0O  in  dissolving  completely 
45  g.  of  Na  or  79  g.  of  K? 

24.  Copy  and  complete.  Do  not  write  in  this  book.  Na  is 
stored  under  ....  When  Na  is  exposed  to  air,  the  formula  of 
the  compound  formed  is  ....  Na  is  used  as  a  catalyst  in  the 
preparation  of  ....  An  instrument  that  makes  use  of  the  ease 
with  which  Cs  loses  electrons  is  the  ....  The  color  imparted 
to  a  flame  by  K  vapor  is  .  .  .,  Li  a  ...  color,  and  Na  a  ... 

25.  (a)  List  four  uses  of  sodium,  and    (b)  three  uses  of 
other  alkali  metals. 

26.  How  is  the  spectroscope  used  in  astronomy? 

27.  Name  three  elements  discovered  by  means  of  the  spec- 

28.  How  can  the  spectrum  of  sunlight  be  obtained? 


29.  What  is  the  difference  in  the  appearance  of  the  spectrum 
of  a   luminous  solid  and   the  spectrum  of  an   incandescent 

30.  Hydrogen  is  included  in  the  replacement  series  of  the 
metals.  Why? 

Group  B 

31.  Zn  appears  higher  than  Fe  in  the  replacement  series. 
What  is  the  reason  for  coating  Fe  with  Zn  to  prevent  corrosion? 

32.  Devise  an  experiment  for  obtaining  Cu  from  CuSO4 

33.  Na  and  K  kept  under  kerosene  for  some  time  lose  their 
silvery  luster.  Explain. 

34.  Discuss  the  use  of  spectroscopy  in  crime  detection. 


1.  Purchase  or  borrow  a  photoelectric  cell  and  connect  it 
in  such  a  way  that,  when  a  flashlight  is  directed  against  it,  a 
bell  will  ring.  Explain  its  action  to  the  class.  What  use  would 
you  make  of  a  photoelectric  cell? 

2.  Make  your  own  replacement  series  of  some  or  all  of  the 
following   metals:    aluminum,   copper,   chromium,   lead,    tin, 
zinc,  and  calcium    (if  you  can  obtain  a  small  piece  from  your 
teacher) .  Use  whatever  suitable  chemicals  you  can  find  around 
the  house  such  as  vinegar.  Report  your  results. 

3.  Prepare   an  illustrated   ten-minute  talk  on  mineral-col- 
lecting as  a  hobby.  Bring  some  of  your  specimens  to  class. 

4.  Write  a  300-500  word  essay  on  one  of  the  various  ana- 
lytical   methods    used    by    chemists    today    such    as     (a)  wet 
method,  (b)  spectrophotometry,  (c)  chromatography,  (d)  tracer 




.  .  .  /  believe  I  speak  for  the  vast  ma- 
jority of  all  scientific  men.  Our  ob- 
ject is  riot  to  make  fobs  and  divi- 
dends. These  are  a  means  to  an  end, 
merely  incidental.  We  wish  to  abol- 
ish drudgery,  discomfort,  and  want 
/row  the  lives  of  men,  and  bring 
them  pleasure,  comfort,  leisure  and 
beauty.  Harold  C.  Urey,  1934 

The  world  receives  a  valuable  gift.  In  1825,  Hans  Christian 
Oersted  (lir'stetfe)  ,  a  Danish  scientist,  announced  that  he  had  isolated 
aluminum  by  gently  heating  aluminum  chloride  and  potassium 
amalgam.  In  1827  Woehler  repeated  Oersted's  experiments  without 
success.  Woehler  finally  obtained  aluminum  by  substituting  metallic 
potassium  for  the  potassium  amalgam  used  by  Oersted. 

One  of  the  first  of  Woehler's  American  students  in  Germany  was 
Professor  Jewett,  of  Oberlin  College.  He  brought  back  to  America 
the  story  of  Woehler's  isolation  of  that  extremely  light,  silvery  metal, 
aluminum,  fewett  was  fond  of  talking  to  his  classes  of  this  strange 
metal,  which  no  one  had  as  yet  been  able  to  obtain  inexpensively  in 
spite  of  its  great  abundance  in  minerals. 

One  day,  as  Professor  Jewett  spoke  of  the  fortune  that  awaited  the 
man  who  was  able  to  develop  a  simple  method  for  extracting  alu- 
minum, one  of  the  students  nudged  his  young  classmate,  Charles 
Martin  Hall.  Chemistry  had  captivated  Hall,  and  his  classmates 
had  known  him  to  make  all  sorts  of  experiments,  hoping  to  make  a 
great  discovery  some  day.  Here  was  his  chance.  His  response  to  that 
nudge  was,  "I  am  going  after  that  metal,"  and  Hall  went  to  work  at 
once  in  his  father's  woodshed 



Hall  attacked  his  problem  scientifically.  He  knew  that  only  the 
most  active  metals,  such  as  sodium  and  potassium,  were  reducing 
agents  strong  enough  to  liberate  aluminum  from  aluminum  chloride. 

A1C13  +  3K  ->  3KC1  +  Al 

Potassium  had  been  used  in  the  method  developed  by  Oersted 
arid  used  later  by  Woehler,  and  sodium  had  been  substituted  tor  po- 
tassium by  the  French  chemist,  Henri  Sainte-Claire  Deville  (saNt- 
klar'  de-vel') .  But  both  potassium  and  sodium  were  too  expensive 
to  use  in  a  commercial  method.  Hall  finally  discarded  all  methods 
that  depended  upon  the  action  of  a  reducing  agent  and  turned  to 
electrolysis,  in  spite  of  his  knowledge  that  Davy,  who  had  isolated 
the  alkali  metals  by  electrolysis,  had  failed  to  get  pure  aluminum  in 
this  way. 

Aluminum  oxide,  called  alumina,  was  the  natural  starting  point. 
Alumina  in  hydrated  form  is  the  chief  component  of  bauxite,  the 
richest  ore  of  aluminum.  But  alumina  has  an  extremely  high  melting 
point.  To  melt  alumina  was  commercially  impracticable.  But  if  an 
electric  current  was  to  liberate  free  aluminum  from  it,  alumina  had 
to  be  either  melted  or  dissolved.  Perhaps  (the  thought  came  to  Hall 
in  one  of  those  flashes  of  genius)  some  mineral  that  would  act  as  a 
solvent  for  aluminum  oxide  might  be  found.  After  trying  a  number 
of  minerals,  he  came  across  a  milky-white,  glassy  solid  called  cryolite. 
He  melted  this  with  some  difficulty  and  then  threw  in  some  alumina. 
The  alumina  dissolved  readily.  He  passed  a  current  through  the 
solution  of  alumina  in  cryolite  and,  to  his  intense  joy,  found  that 
metallic  aluminum  was  deposited  at  the  cathode 

On  February  23,  1886,  he  burst  excitedly  into  the  laboratory  of 
Professor  Jewett  and,  holding  a  few  aluminum  "buttons,"  exclaimed, 


Charles  Martin  Hall,  the  discoverer  of  the 
electrolytic    process    of    producing   alumi- 

p,     «o    c  ,-        ,  +       I.     .1.     4. ^busbar 

Fig.    89.   Cross   section    of   a  '..        \m\      i,i      D        ..'    ^,^Mt«»  JL  A!  o 

*  ,      ..  .    .     .        .  M      TT     T      •        rTx^'^crusF  or  AUw3 

furnace  in  which  aluminum  is  I]         1        _  |  _     I    ^Jr[ 

produced  commercially  by  the  ^j^jjjAjjjLj^^          AI2O3  In 

Hall  process.       carbon  anode  +  ^g^^^^^J^Tn  fused^cryolite 

carbon  lining  cathode  aluminum 

"I've  got  it!"  Hall  soon  obtained  a  patent  on  his  process  and  two 
years  later  the  Pittsburgh  Reduction  Company,  which  later  became 
the  Aluminum  Company  of  America,  was  founded.  In  1914  Hall 
died,  world-famous  and  a  multimillionaire.  He  left  most  of  his  for- 
tune to  Oberlin  College  and  other  educational  institutions. 

Discoveries  often  result  from  social  needs.  Hall  was  22  when  he 
produced  aluminum.  Exactly  two  months  after  Hall  had  handed  his 
teacher  the  first  samples  of  his  aluminum,  another  chemist,  Paul 
Heroult  (a-roolt') ,  also  22  years  of  age,  applied  for  a  patent  in  France 
on  an  identical  commercial  method  of  preparing  aluminum. 

This  is  not  a  rare  example  of  simultaneous  discoveries  in  the  his- 
tory of  science.  Advances  in  science  often  are  made  in  different  parts 
of  the  world  at  almost  the  same  time.  They  are  frequently  the  final 
step  in  a  long  series  of  experiments  conducted  by  many  research 
workers  in  widely  separated  laboratories.  The  scientist  who  is  fortu- 
nate enough  to  publish  his  discovery  first  is  recognized  as  the  hero 
of  a  battle  in  which  many  other  soldiers  have  been  engaged.  The 
heroes  of  science,  on  the  whole,  concede  this  element  of  good  for- 
tune. Can  you  think  of  another  such  instance  in  the  story  of  scientific 

Metallurgy  of  aluminum  by  the  Hall  process.  Hall's  process  is  still 
in  use.  The  electric  furnace  used  is  an  open  cell  about  25  to  30  feet 
long,  lined  with  carbon,  which  constitutes  the  cathode.  Powdered 
cryolite  mined  in  Greenland  or  made  synthetically  is  placed  in  the 
cell,  and  as  the  current  passes  through  it,  it  melts.  Aluminium  oxide, 
or  alumina,  is  a  white  powder  obtained  by  refining  bauxite  ore.  It 
is  added  to  the  molten  cryolite  and  immediately  dissolves.  The 
aluminum  oxide  dissociates,  forming  positive  aluminum  ions  and 
negative  oxygen  ions. 

A12O3  -»  2A1+++  +  3O— 

Carbon  rods  are  suspended  in  the  molten  aluminum  oxide  solution 
and  act  as  the  anode.  When  the  circuit  is  closed,  the  aluminum  ions 
travel  to  the  cathode,  where  they  obtain  electrons  which  change 
them  from  aluminum  ions  to  free  aluminum.  This  free  molten  alu- 
minum then  settles  to  the  bottom  of  the  cell.  Later,  a  hole  at  the 



bottom  of  the  cell  is  unplugged,  and  the  molten  aluminum  is  tapped 
off  into  large  ladles  and  cast  in  molds,  in  which  it  solidifies  as  pig 
aluminum.  The  oxygen  ions,  in  the  meantime,  have  traveled  to  the 
anode,  given  up  their  electrons,  and  changed  to  free  oxygen.  This 
oxygen  combines  with  the  carbon  of  the  anode  and  forms  carbon 
dioxide,  given  off  as  a  gas. 

c  -f  o2  -» co2  T 

The  process  is  continuous.  Aluminum  oxide  is  added,  aluminum  is 
removed,  and  the  carbon  anodes  are  replaced  from  time  to  time. 
The  original  cryolite,  Na3AlF«,  though  it  contains  aluminum,  does 
not  decompose.  It  acts  only  as  a  solvent.  Many  of  these  electrolytic 
cells  are  joined  in  series.  Commercial  cells  produce  about  500  pounds 
of  99+  percent  pure  aluminum  per  day. 

Bauxite,  found  in  large  amounts  in  Surinam  (Dutch  Guiana)  , 
and  British  Guiana  —  Arkansas  leads  the  United  States  in  the  pro- 
duction of  this  ore  —  contains  a  fair  percentage  of  the  oxides  of  iron, 
silicon,  and  titanium.  If  these  impurities  are  not  removed  before 
the  bauxite  is  added  in  the  electric  furnace,  the  aluminum  produced 
is  impure. 

After  the  aluminum  is  drawn  from  the  electrolytic  cells,  the  pig 
metal  is  remelted  so  that  the  nonmetallic  impurities  may  be  skimmed 
off.  If  aluminum  alloys,  rather  than  pure  aluminum,  are  desired,  the 
alloying  may  be  done  during  the  remelting.  The  chief  alloying  ele- 
ments include  copper,  magnesium,  manganese,  silicon,  zinc,  iron, 
nickel,  and  chromium. 

The  physical  properties  of  aluminum.  Aluminum  is  silvery  white 
in  color  and  is  one  of  the  lightest  of  the  common  metals.  It  is  only 
one-third  as  heavy  as  iron.  It  is  very  malleable  and  ductile  and  com- 
pares well  with  both  silver  and  copper  in  the  ease  with  which  it  con- 
ducts both  heat  and  electricity.  It  can  be  worked  readily;  that  is,  it 
can  be  cast,  rolled,  forged,  extruded,  machined,  or  drawn.  Parts  can 
be  joined  by  welding,  brazing,  and  riveting. 

The  chemical  properties  of  aluminum.  The  atomic  weight  of  alu- 
mium  is  27.  Its  atomic  number  is  13;  hence  it  has  only  three  elec- 
trons in  its  outside  ring.  It  is,  therefore,  a  metal  with  a  valence  of 
plus  three.  Aluminum  is  an  amphoteric  element  and  may  act  as  either 
an  acid  or  a  base.  It  is  attacked  by  strong  bases  as  follows: 

2A1  +  2NaOH  +  2H2O  -»  3H2 1  +  2NaAlO2  (sodium  aluminate) 

This  sodium  aluminate  is  the  salt  of  aluminic  acid,  H3A1O3.  Be- 
cause of  the  reaction  between  aluminum  and  the  strong  bases,  or 



substances  with  basic  reactions,  such  as  washing  soda,  care  should  be 
taken  not  to  heat  such  substances  in  aluminumware. 
Aluminum  is  attacked  by  nearly  all  acids,  forming  aluminum  salts. 

2A1  +  6HC1  -»  3H2 1  +  2A1C13 

The  surface  of  aluminum  oxidizes  rapidly  in  air,  forming  alu- 
minum oxide,  AL,O3.  This  extremely  thin,  transparent,  but  tough 
film  acts  as  an  excellent  protective  coating  and,  unlike  iron  rust, 
adheres  firmly  to  the  surface  of  the  metal,  thus  preventing  further 
oxidation  unless  the  coating  is  perforated.  Alclad  is  a  sheet  of  alu- 
minum alloy  such  as  duralumin  covered  with  a  layer  of  aluminum. 
It  resists  corrosion  very  well. 

Tremendous  growth  of  the  aluminum  industry.  Before  the  Hall 
process  was  introduced,  aluminum  was  not  used  widely,  because  of 
the  great  cost  of  preparing  it.  It  is  a  far  cry  from  the  world  produc- 
tion of  two  tons  of  aluminum  in  1859  at  17  dollars  per  pound  to  the 
more  than  one  million  tons  of  this  metal  produced  in  the  United 
States  and  Canada  alone  in  a  recent  year  at  about  20  cents  a  pound. 

The  Hall  process  gives  primary  aluminum,  that  is,  metal  pro- 
duced directly  from  an  ore  or  ores.  But  a  significant  and  perhaps 
increasing  source  of  aluminum,  and  other  metals  as  well,  lies  in 
secondary  sources,,  that  is,  sources  from  which  a  metal  is  recoverable 
from  one  use  for  reuse  in  another.  Secondary  or  scrap  aluminum  is 
a  very  large  source  of  pure  aluminum.  Production  of  metals  from 

Surface  mining  of  bauxite  in  Surinam.  Over-lying  earth  is  first  removed.  The  ex- 
posed ore  is  then  loosened  by  blasting  and  loaded  into  the  mine  cars. 

Aluminum  Company  of  America 

An  "aluminum  skyscraper"  in  Pitts- 
burgh, Pa.  This  30-story  building  is 
constructed  of  aluminum  panels 
mounted  on  a  steel  framework.  The 
ceilings,  wiring,  ventilation  ducts, 
doors,  hardware  and  most  of  the 
plumbing  are  made  of  aluminum. 

secondary  sources  is  one  phase  of  an  intelligent  metals  conservation 
program.  Such  a  program  can  do  much  to  conserve  natural  resources. 

By  far  the  largest  users  of  aluminum  are  the  transportation  indus- 
tries. Great  quantities  of  the  "metal  with  wings"  are  used  in  the 
construction  of  airplanes,  streetcars,  railroad  cars,  locomotives,  steam- 
ships, motorships,  automobiles,  trucks,  buses,  bicycles,  and  motor- 

Through  decreasing  the  weight  of  such  carriers  as  airplanes,  rail- 
road cars,  and  trucks,  payload  and,  thereby,  revenue  can  be  increased. 
At  the  same  time,  if  the  reduction  in  weight  applies  only  to  the  total 
gross  weight  of  the  vehicle,  as  in  an  automobile  or  bus,  much  less 
energy  is  required  to  attain  and  maintain  speed,  much  less  energy 
is  lost  in  stopping,  and  operating  costs  are  thereby  reduced. 

Electricity  and  aluminum.  The  electric  industries  use  thousands 
of  tons  of  aluminum  yearly  in  lines  for  the  transmission  of  electricity 
over  long  distances.  For  this  purpose,  aluminum  cable  with  a  steel 
reinforcing  core  is  used  instead  of  copper.  Since  aluminum  is  lighter 
than  copper,  fewer  towers  are  required  to  support  the  cables.  More 
than  1.5  million  miles  of  aluminum  transmission  cables  carry  elec- 
tricity to  almost  all  parts  of  this  nation.  Aluminum  is  used  also  in  the 
production  of  parts  for  electric  equipment  of  many  kinds,  such  as 



vacuum  cleaners  and  various  household  appliances,  and  particu- 
larly in  parts  for  radios  and  other  electronic  equipment. 

Aluminum  in  the  kitchen.  One  use  of  aluminum  goes  back  to 
1890,  when  the  first  aluminum  cooking  utensils  were  produced. 
Since  aluminum  is  an  excellent  conductor  of  heat,  and  at  the  same 
time  is  very  light,  aluminum  cooking  utensils  are  very  popular. 

Certain  alkaline  foods  and  waters  heated  in  aluminum  may  pro- 
duce a  superficial  discoloration,  which  is  harmless  and  readily  re- 
moved by  a  mild  abrasive  cleaner.  Do  not  permit  all  the  cooking 
liquid  in  a  lightweight  aluminum  utensil  to  boil  off,  for  if  this 
occurs,  a  hole  may  be  "burned"  in  the  bottom  of  the  utensil  as  a 
result  of  the  relatively  low  melting  point  of  aluminum. 

Because  strong  alkalies  attack  the  protective  coating  of  alumi- 
num oxide  that  forms  on  the  surface  of  aluminum  as  well  as  the 
aluminum  itself,  cooking  utensils  made  of  the  metal  should  not  be 
scrubbed  or  polished  with  harsh  alkali  cleaners.  To  clean  aluminum 
utensils  use  soap  and  water  or  mild  abrasive  cleansers  only. 

Other  uses  of  aluminum.  In  the  packaging  of  foods  and  other  com- 
modities, aluminum  foil  has  almost  entirely  replaced  tin  foil.  Candy 
bars,  chewing  gum,  cream  cheese,  camera  film,  and  countless  other 
articles  go  to  market  in  shining  dress.  Aluminum  foil  coated  with  a 
plastic  film  is  suitable  for  the  packaging  of  almost  any  kind  of  food. 
Aluminum  leaf  is  used  in  photoflash  lamps. 

Collapsible  tubes  made  of  aluminum  carry  such  items  as  shaving 
cream,  tooth  paste,  and  cosmetics,  while  vital  serums  and  various 
other  pharmaceutical  preparations  are  packed  in  glass  bottles  with 
aluminum  seals.  Aluminum  paints  are  widely  used  for  protecting 
both  wood  and  metal.  Aluminum  foil  is  used  for  home  insulation. 

Aluminum  Company  of  America 

Circular  aluminum  blanks 
being  removed  from  the 
conveyor  of  an  annealing 
furnace.  These  blanks  will  be 
used  in  the  manufacture  of 
cooking  utensils. 





Much  furniture  and  many  decorative  articles  for  the  household  are 
made  of  aluminum. 

The  buildings  of  Rockefeller  Center  in  New  York  contain  more 
than  1000  tons  of  aluminum  in  the  vertical  panels  between  windows. 
In  the  finishing  of  steel,  large  quantities  of  aluminum  are  used  in 
removing  oxides  from  the  molten  steel.  The  largest  use  of  aluminum 
is  in  the  form  of  alloys  of  much  greater  strength  than  pure  aluminum. 

Thermit  is  used  in  welding.  Attempts  to  reduce  aluminum  oxide 
with  carbon  failed  because  of  the  great  attraction  of  aluminum  for 
oxygen.  Aluminum  is  a  powerful  reducing  agent,  especially  when 
it  is  in  the  form  of  a  fine  powder.  Because  a  powder  has  a  much 
greater  reacting  surface  than  a  solid,  a  powder  makes  possible  a  more 
intense  chemical  reaction  than  the  same  weight  of  the  same  solid  in 
larger  pieces. 

When  a  mixture  of  powdered  aluminum  and  iron  oxide  is  ignited 
by  means  of  a  fuse,  such  as  a  strip  of  magnesium  ribbon,  a  chemical 
reaction  takes  place  at  once  in  which  the  aluminum  takes  the  oxygen 
away  from  the  iron  oxide,  leaving  a  residue  of  pure  iron. 

8A1  +  3Fe3O4  -»  4A12O3  +  9Fe 

The  heat  of  this  reaction  is  so  great  that  the  iron  formed  is  molten. 
This  mixture  of  aluminum  and  iron  oxide,  known  as  thermit,  is  used 
in  welding  broken  propeller  shafts,  rudder  frames,  locomotive  parts, 
and  in  situations  where  repairs  must  be  made  on  the  spot.  It  is  also 
used  in  one  type  of  incendiary  bomb. 

The  metals  chromium  and  manganese  (and  wolfram,  vanadium, 
molybdenum,  silicon  and  boron)  may  be  extracted  from  their  oxides 
or  ores  by  aluminothermy,  an  aluminum  reduction  similar  to  the 
thermit  reaction.  The  equations  for  the  reduction  of  chromium  and 
manganese  ore  are: 

Cr2O3  +  2A1  ->  A12O3  +  2Cr 
3Mn3O4  +  8A1  -»  4A12O3  +  9Mn 

This  is  a  common  way  of  manufacturing  or  producing  such  metals. 
What  are  alums?  When  potassium  sulfate  is  dissolved  in  a  solution 
of  aluminum  sulfate,  the  two  salts  combine  and  crystallize  out  as 
potassium  aluminum  sulfate. 

K2SO4  +  A12(SO4)3  +  24H2O  ->  2KA1(SO4)2     12H2O 

This  compound  is  called  common  alum.  It  is  one  of  a  group  of  salts, 
called  the  alums,  which  resemble  one  another  in  the  eight-sided  form 
of  their  crystals,  in  their  solubility  in  water,  and  in  their  type  for- 
mula. The  type  formula  of  an  alum  is  XY  (SO4)  L>  •  12H2O.  In  it  X  may 
be  K,  Na,  or  NH4,  and  Y  may  be  a  trivalent  element,  such  as  Al,  Fe, 
or  Cr.  An  alum  is  a  double  salt,  that  is,  a  salt  containing  two  metals 
and  one  acid  radical. 

Common  alum  has  a  sweetish  taste  and  is  used  in  "Foamite,"  mor- 
danting (see  page  605)  ,  and  water  purification.  Alum  is  used  also 
in  making  alum  baking  powder  and  as  an  astringent,  a  substance  that 
contracts  skin  tissues.  It  is  used  in  the  sizing  of  paper.  Chrome  alum 
is  used  in  the  tanning  of  leather,  during  which  skins  or  hides  are 
made  softer,  and  more  resistant  to  the  action  of  bacteria. 

Aluminum  hydroxide,  its  preparation  and  properties.  When  any 
soluble  aluminum  salt  is  added  to  water  or  to  ammonium  hydroxide, 
aluminum  hydroxide,  a  white,  gelatinous  precipitate,  may  be  formed. 
The  aluminum  salt  is  said  to  hydrolyze,  since,  as  in  the  case  of  the 
solution  of  aluminum  chloride  in  water,  an  acid  and  a  base  are 

A12(SO4)3  +  6NH4OH  • 
A1C13  +  3H2O 

-  2A1(OH)3 1  +  3(NH4)2SO4 

•  3HC1  +  A1(OH)8 1 


Use  by  approximate 

Industrial,  agricultural, 
and  mining  machinery 


Aluminum  hydroxide  may  act  either  as  an  acid  or  as  a  base.  When 
it  comes  in  contact  with  a  base,  it  acts  as  a  weak  acid  and  combines 
with  the  base. 

HaAlOs  +  NaOH  ->  NaAlO2  (sodium  aluminate)  +  2H2O 

A1C18  is  the  active  constituent,  about  15  percent,  of  many  body 
deodorants  and  antiperspirants. 

Aluminum  sulfate  in  the  purification  of  water.  As  you  know, 
aluminum  sulfate,  when  it  is  added  to  water,  hydrolyzes,  forming 
jelly-like  aluminum  hydroxide. 

A12(S04)3  +  6H20  -»  2A1(OH)3  j  +  3H2SO4 

This  aluminum  hydroxide  gradually  settles  to  the  bottom  and  carries 
down  with  it  any  particles  that  are  floating  in  the  water,  including 
bacteria,  industrial  wastes,  and  fine  clay.  This  process,  used  in  the 
purification  of  water,  is  called  coagulation  and  is  the  first  step  in 
clearing  water  of  its  turbidity.  Coagulation  does  not  remove  dis- 
solved impurities. 

Ferrous  sulfate  is  sometimes  used  instead  of  aluminum  sulfate, 
because  it  also  forms  a  gelatinous  precipitate,  ferrous  hydroxide,  in 
water.  But  ferrous  hydroxide  rapidly  oxidizes  to  ferric  hydroxide, 
which  is  also  gelatinous.  Hence,  it  is  the  ferric  hydroxide  that  actu- 
ally reduces  the  turbidity,  or  clarifies  the  water. 

FeS04  +  2H20  ->  H2SO4  +  Fe(OH)2  j    (ferrous  hydroxide) 

Aluminum  oxide  — ore,  gem,  abrasive,  refractory.  When  alumi- 
num hydroxide  is  heated,  it  forms  a  white,  insoluble  oxide  of  alumi- 
num, which  melts  above  3600° F. 

2A1(OH)3  ->  A12O3  +  3H2O 

As  hydrated  oxides,  A12O3  •  3H2O  and  A12O8  •  H2O  are  found  widely 
distributed  in  the  aluminum  ore,  bauxite.  The  precious  gem  stones 
ruby  and  sapphire  are  composed  of  alumina,  colored  by  the  pres- 
ence of  small  amounts  of  metal  oxides.  Successful  methods  have 
been  developed  to  prepare  synthetic  rubies  and  sapphires  by  melting 
pure  aluminum  oxide  in  the  heat  of  an  oxy hydrogen  flame.  These 
synthetic  stones  cannot  easily  be  distinguished  from  natural  gems. 

The  production  of  synthetic  rubies  and  sapphires  is  increasingly 
important,  for  they  are  used  as  bearings  in  watches,  electric  indica- 
tors, sensitive  electric  relays,  and  in  thousands  of  other  kinds  of  pre- 
cision instruments. 


Emery  is  a  natural  aluminum  oxide,  which  is  extremely  hard  and 
can  be  used  as  an  abrasive  for  grinding,  polishing,  drilling,  and  cut- 
ting. It  is  almost  as  hard  as  diamond.  Fused  alumina,  for  use  as  an 
abrasive,  is  prepared  in  large  quantities  by  the  fusion  of  alumina  in 
an  electric  furnace.  One  trademark  for  such  fused  alumina  is  "Alun- 

Not  only  is  fused  alumina  a  good  abrasive,  but  because  of  its  high 
melting  point,  it  is  an  excellent  refractory.  A  refractory  is  a  sub- 
stance which,  because  it  melts  at  a  high  temperature  and  is  chem- 
ically inert,  can  be  used  for  furnace  linings  and  in  similar  installa- 
tions which  must  withstand  extreme  heat.  Fused  alumina  refractories 
are  therefore  used  in  making  bricks,  spark  plugs,  crucibles,  cements 
for  high-temperature  work,  and  high-temperature  thermometers 
called  pyrometers.  Because  alumina  is  quite  inert  chemically,  it  is 
employed  as  a  catalyst  support,  and  for  some  laboratory  ware  such 
as  porous  plates  used  in  filtering  chemical  solutions. 

Activated  alumina  is  a  highly  porous  A12O3  which  is  used  to  adsorb 
moisture  in  air  conditioning,  and  to  dry  such  gases  as  propane  and 

Newcomers,  the  light  alloys  of  aluminum.  The  American  Society 
for  Testing  Materials  defines  an  alloy  as:  A  substance  consisting  of 
two  or  more  metallic  elements  or  of  metallic  and  nonmetallic  ele- 
ments which  are  soluble  in  each  other  when  molten,  and  which  do 
not  separate  into  distinct  layers  when  solid.  The  art  of  making  alloys 
is  very  old.  One  of  the  oldest  alloys  is  bronze.  This  alloy  gave  its 
name  to  an  era  of  human  progress,  the  Bronze  Age,  which  began 
about  2500  B.C.,  and  gave  way  to  the  Iron  Age  some  15  to  20  cen- 
turies later. 

Alloys  are  in  many  cases  neither  true  solutions  nor  compounds. 
Whereas,  for  example,  Cu  and  Au  mix  in  all  proportions,  others 
have  a  limited  solubility.  Some  alloys  consist  of  metals  present  in 
almost  constant  proportions  such  as  NaZnxl.  The  properties  of  an 
alloy  may  differ  radically  from  those  of  the  elements  that  compose 
it.  An  alloy  is  usually  harder  and  more  resistant  than  any  of  its  con- 
stituents. For  these  reasons,  alloys  are  extremely  important.  By  mix- 
ing two  or  more  metals,  we  may  obtain  an  alloy  whose  properties 
are  immensely  more  valuable  than  those  of  any  of  its  ingredients. 
For  example,  the  tensile  strength  of  many  aluminum  alloys  is  greater 
than  that  of  pure  aluminum. 

Some  of  the  alloys  of  aluminum  containing  copper,  manganese, 
magnesium,  or  other  metals  are  light,  strong,  and  easily  machined. 
Such  alloys  make  ideal  materials  for  the  framework  of  the  wings  and 





fuselages,  and  landing  gear  and  propellers  of  airplanes,  and  for  cer- 
tain parts  of  automobiles  and  railway  cars.  "Duralumin,"  was  one  of 
the  first  strong  aluminum  alloys  to  be  given  additional  strength  by 
heat  treatment.  Other  alloys,  stronger,  tougher,  and  more  durable, 
have  largely  replaced  it. 

Some  of  the  modern  alloys  of  aluminum  are  actually  stronger  than 
structural  steel,  and  much  lighter.  Manufacturers  of  aluminum  alloys 
call  their  products  by  many  different  names.  Aluminum  alloys  of  the 
Aluminum  Company  of  America  are  called  "Alcoa"  alloys,  and  usu- 
ally a  number  indicates  the  particular  "Alcoa"  alloy. 

Whereas  there  are  about  40  metals  in  common  use,  more  than 
6000  alloys  have  helped  make  possible  our  modern  industrial  world. 
If  we  try  to  escape  from  alloys,  we  escape  from  civilization.  Steel, 
the  most  important  single  alloy  of  all,  is  discussed  in  Chapter  27. 

Magnesium,  lightweight  of  the  structural  metals.  Magnesium  is 
one  of  the  metals  most  commonly  alloyed  with  aluminum.  It  can  also 
be  used  for  many  of  the  same  purposes  as  aluminum  with  which  it 
has  become  competitive. 

World  War  II,  with  its  tremendous  military  demand  for  light  met- 
als, saw  the  infant  magnesium  industry  grow  from  a  healthy  baby  to 
a  full-grown  adult  almost  overnight.  Production  of  magnesium 
jumped  in  the  five  years  following  1939  from  3500  tons  to  a  high  of 
nearly  200,000  tons.  Volume  for  volume,  magnesium  in  1953  at  26.5 
cents  a  pound  was  less  expensive  than  aluminum  selling  at  20.9  cents 
a  pound. 

Magnesium  is  a  soft,  silvery-white  metal  about  60  percent  as  heavy 
as  aluminum.  In  pure  form  it  has  little  structural  strength,  but 
when  properly  alloyed,  the  resulting  materials  have  good  structural 
strength.  Although  the  strengths  of  magnesium  alloys  are  less  than 
those  of  the  heavier  structural  metals,  by  using  slightly  greater  thick- 
nesses of  the  alloys,  structural  shapes  as  strong  as  steel,  but  with  only 
a  fraction  of  the  weight  of  steel,  may  be  made. 

"Dowmetal"  is  the  trademark  of  a  family  of  magnesium  alloys 
made  by  the  Dow  Chemical  Company,  pioneer  producers  of  magne- 
sium. All  "Dowmetals"  are  composed  of  magnesium  and  varying 
percentages  of  other  metals  such  as  manganese,  aluminum,  or  zinc. 
The  percentages  of  these  metals  in  the  alloys  vary  with  the  char- 
acteristics desired  in  the  alloy.  For  example,  "Dowmetal  A"  contains 
eight  percent  aluminum  and  92  percent  magnesium.  "Dowmetal  G" 
contains  ten  percent  aluminum,  0.1  percent  manganese,  and  89.9 
percent  magnesium.  A  similar  series  of  magnesium  alloys  is  manu- 
factured by  the  American  Magnesium  Corporation.  Such  alloys  are 
designated  by  number,  for  example,  "AM3S"  or  "AM57S."  Magne- 
lite,  used  in  kitchenware,  contains  Al,  Mg,  and  silicon. 

Chemical  properties  of  magnesium.  Magnesium  is  high  in  the 
replacement  series  of  metals,  but  pure  magnesium  and  some  of  its 
alloys  corrode  relatively  slowly  even  in  moist  air.  Like  aluminum, 
magnesium  is  a  self-protecting  metal.  The  protective  film  is  probably 
a  basic  carbonate.  Powdered  magnesium  or  thin  strips  of  magnesium 
burn  in  air  when  ignited,  forming  magnesium  oxide.  Hot  magne- 
sium burns  in  CO2  to  form  MgO  and  C. 

2Mg  +  C02  ->  2MgO  +  C 

Fused  magnesium  oxide  under  the  name  of  "magnorite"  is  a  refrac- 

Magnesium  from  the  sea.  Magnesium  is  one  of  the  most  abundant 
industrial  metals  in  the  earth's  crust.  There  is  an  almost  unlimited 
supply  of  it  in  the  sea,  for  one  cubic  mile  of  sea  water  contains  nine 
billion  pounds  of  this  metal  in  the  form  of  MgCl2.  The  brine  wells, 
such  as  those  near  Midland,  Michigan,  are  another  source  of  magne- 
sium chloride.  Until  rather  recently,  magnesium  production  in  the 
United  States  was  limited  almost  completely  to  the  electrolysis  of 
molten  magnesium  chloride  obtained  from  the  sea  or  salt  wells,  or 
of  fused  carnallite,  MgCL  •  KC1.  During  electrolysis,  the  magnesium 

These  workmen  are  pouring 
molten  magnesium  into  a 
mold.  Magnesium  can  be 
cast,  molded,  extruded,  and 
worked  by  many  of  the  com- 
mon methods  of  the  metals 

collects  at  the  cathode,  and  chlorine  at  the  anode.  Magnesium  is  thus 
the  first  structural  metal  obtained  from  the  sea. 

Another  source  of  magnesium  is  dolomite,  MgCO;,  •  CaCO,,  which 
is  very  widely  distributed  throughout  the  world,  forming  entire 
mountain  ranges.  Magnesium  has  been  obtained  from  this  ore  by 
roasting  it  to  form  MgO,  and  then  reducing  the  magnesium  oxide 
at  high  temperatures  with  carbon. 


Holmes,  Harry  N.  Out  of  the  Test  Tube.  Emerson  Books, 
New  York,  1945.  "Three  Light  Metals"  are  discussed  in  this 
very  popular  science  book. 

Forbes,  R.  J.  Man,  the  Maker.  Harry  Schuman,  New  York, 
1950.  An  excellent  and  compact  history  of  technology  and 
engineering  including  a  short  account  of  the  Hall  process. 

Wade,  Frank  B.  "Man-Made  Gems."  Journal  of  Chemical 
Education,  June,  1931,  pp.  1015-1026.  Describes  the  manufac- 
ture of  artificial  rubies  and  sapphires,  and  the  methods  used  in 
detecting  differences  between  natural  and  artificial  gems. 


1.  An  alloy  is  a  substance  consisting  of  two  or  more  metallic 
elements  or  of  metallic  and  nonmetallic  elements  that  are 
soluble  in  each  other  when  molten  and  do  not  separate  into 
distinct  layers  when  solid. 



2.  A  double  salt  is  a  salt  that  contains  two  metals  and  one 
acid  radical. 

3.  Frequently,  scientific  discoveries  and  inventions  are  made 
independently  and  simultaneously  in  different  laboratories. 


Group  A 

1.  By  whom  was  Al  isolated  in  1825? 

2.  (a)  When  and  by  whom  was  the  first  successful  com- 
mercial method  of  producing  Al  devised?   (b)  What  were  the 
circumstances  surrounding  his  solution  of  this  problem? 

3.  Name  five  elements  that  are  prepared  by  electrolysis. 

4.  Write    an    equation    showing   how   Woehler   obtained 
pure  Al. 

5.  In  the  metallurgy  of  Fe,  C  is  used  to  reduce  Fe2O3.  Why 
is  C  not  used  to  liberate  Al  from  its  oxide? 

6.  (a)    Make  a  labeled  diagram  of  the  electrolytic  cell  in 
which  Al  is  freed  from  its  oxide,    (b)  Discuss  in  detail  the 
chemistry  of  the  extraction  of  Al  by  the  Hall  process. 

7.  (a)  What  is  the  function  of  Na3AlF0  in  the  Hall  process? 
(b)  What  evidence  is  there  that  the  Na3AlF6  does  not  supply 
some  of  the  Al  formed  at  the  cathode? 

8.  (a)   Is  bauxite  just  as  it  comes  from  the  mines  used  in 
extracting  Al  by  the  Hall  process?  (b)  Explain. 

9.  Give  reasons  why  each  of  the  following  statements  is 
either  true  or  false:    (a)  O2  is  liberated  at  the  anode  of  the 
aluminum  cell,   (b)  The  carbon  anodes  of  the  aluminum  cell 
must  be  replaced  from  time  to  time,  (c)  Much  Al  is  produced 
from  alumina  in  plants  located  near  Niagara  Falls. 

10.  Find  the  percentage  of  Al  in    (a)  pure  alumina  and 
(b)  in  pure  Na3AlF0. 

11.  Aluminum   hydroxide    may   be    written   Al  (OH)  3    or 
H3A1O3.  Explain. 

12.  Lustrous  Al  becomes  covered  with  a  coating  in  air.  Ex- 
plain this  chemical  change. 

13.  Al  is  higher  in  the  replacement  series  of  the  metals  than 
Fe,  but  it  corrodes  much  less  than  Fe.  Explain. 


14.  Write   the  equations  for:    (a)  Action  of  hydrochloric 
acid  on  aluminum,   (b)  Action  of  a  strong  base  on  aluminum, 
(c)  What  precautions  are  necessary  in  using  aluminumware 
because  of  this  action  of  bases? 

15.  300  Ib.  of  powdered  Fe3O4  are  available.  How  much 
thermit  mixture  can  be  prepared  from  this  Fe3O4? 

16.  (a)  What  properties  of  Al  make  it  a  satisfactory  material 
for  use  in  construction  work?    (b)  Why  cannot  Al  displace 
steel  completely  in  construction  work? 

17.  (a)  What  are  the  most  important  uses  of  Al?    (b)  Op- 
posite each  use  write  the  properties  that  make  this  use  pos- 

_  t .  . . 

18.  Make  a  list  of  ten  articles  used  by  you  every  day  which 
contain  Al. 

19.  (a)  What  is  an  alloy?  (b)  Give  three  examples. 

20.  What  are  the  composition  and  uses  of  one  of  the  chief 
alloys  of  Al?  i 

21.  How  is  a  photoflash  bulb  constructed? 

22.  (a)  Of  what  is  thermit  composed?    (b)  For  what  is  it 
used?  (c)  Write  the  equation  for  the  reaction  that  takes  place 
when  thermit  is  ignited,    (d)  What  is  the  source  of  the  heat 
produced?   (The  temperature  is  so  high  that  Fe,  which  melts 
at  1530°C.,  becomes  liquid.) 

23.  (a)  What  two  metals  are  extracted  from  their  ores  by 
the  use  of  Al?  (b)  Explain,  using  an  equation. 

24.  (a)  What  is  a  double  salt?  (b)  Give  one  illustration. 

25.  Explain  how  it  is  possible  to  have  an  alum  that  does  not 
contain  Al. 

26.  Write  an  equation  showing  how  aluminum  hydroxide 
is  prepared  in  the  laboratory. 

27.  (a)  What  property  of  Al  (OH)  3  makes  it  useful  in  puri- 
fying water?   (b)  Of  what  kind  of  impurities  does  it  free  the 
water?    (c)  How?    (d)  What  is  added  to  the  water  to  form 
the  Al  (OH)  3? 

28.  Explain  how  artificial  rubies  are  prepared. 

29.  (a)  How  is  fused  alumina  made?  (b)  What  is  activated 
alumina?  (c)  What  are  some  of  its  uses? 

30.  (a)  Describe  two  chemical  properties  of  Mg.    (b)  For 
what  is  Mg  used? 

31.  (a)  Wfcat  metal  is  the  chief  competitor  of  Al?  (b)  Why? 


Group  B 

32.  (a)  What  is  the  effect  of  A12  (SO4)  3  on  moist  blue  litmus 
paper?   (b)  Explain. 

33.  Describe  one  industrial  process  for  the  production  of  Mg. 

34.  (a)  What  is  a  refractory?    (b)  Name  two  refractories 
that  are  used  commercially  in  large  quantities. 

35.  Mg  is  a  very  abundant  metal.  Why  has  it  only  lately 
come  into  widespread  use? 

36.  One  cu.  ft.  of  pure  water  weighs  62.5  Ib.  What  does  one 
cu.  ft.  of  pure  Al  weigh? 

37.  Washing  soda  should  not  be  boiled  in  aluminum  uten- 
sils. Why? 

38.  Name  three  chemical  discoveries  that  were  made  by 
young  men.  Do  not  include  Hall's  discovery. 

39.  Compare  the  properties  of  an  alloy  with  the  properties 
of  (a)  a  compound,  (b)  a  solution,  and  (c)  a  mixture. 

40.  (a)  What  kind  of  salts  hydrolyze  in  H2O?  (b)  Explain. 

41.  Al  is  the  most  abundant  metal  in  the  earth's  crust.  How- 
ever, Fe  costs  less  than  Al.  Explain. 

42.  (a)  Write  the  reaction  that  occurs  when  A12  (SO4)  3  is 
placed  in  water,    (b)  Write   the  reaction  that  occurs  when 
Na2CO8  is  placed  in  water,  (c)  What  ions  exist  in  each  of  these 
solutions?   Why?    (d)  If  we  add  A12  (SO4)  s  to  a  solution  of 
Na2CO3,  we  get  Al  (OH)  3  and  not  A12  (CO3)  3.  Why? 

43.  The  thermit  process  has  been  called  a  vestpocket  blast 
furnace.  Explain  this  comparison. 


1.  Make  a  small  model  of  the  electrolytic  cell  used  in  mak- 
ing pure  aluminum. 

2.  Make  a  collection  of  articles  or  samples  of  alloys  contain- 
ing aluminum.  Report  to  the  class  on  the  composition  and  use 
of  each  item  collected. 

3.  At  the  hundredth  anniversary  of  the  opening  of  the  U.S. 
Patent  Office   (1936)   the  12  greatest  inventions  made  in  this 
country  were  listed.  Low-cost  aluminum  was  included.  Make  a 
list  of  11  other  inventions  which  you  might  have  included  in 
such  a  list.  Consult  your  teacher  of  history. 




To  an  extent  not  generally  appreci- 
ated, U.S.  industrial  strength  is  based 
on  an  accident  of  nature:  the  unique 
geographical  combination  of  Minne- 
sota and  Michigan  ore,  Appalachian 
coking  coal  and  the  Great  Lakes 
highway.  On  this  triple  gift  of  nature 
rests  our  towering  steel  industry.  To 
it  we  owe,  in  the  final  analysis,  our 
standard  of  living.  Leonard  Engel  in 
Scientific  American,  May,  1948 

Importance  of  steel  in  our  machine  age.  In  a  recent  year,  more 
than  100  million  tons  of  steel  were  produced  to  satisfy  the  appetite 
of  the  machine  age.  The  iron  produced  in  an  entire  year  a  hundred 
years  ago  would  meet  our  present  needs  for  just  one  day.  This 
fact  is  easy  to  understand  when  we  think  of  the  great  skyscrapers, 
the  thousands  of  miles  of  rails,  the  millions  of  automobiles,  the 
thousands  of  locomotives,  ships,  and  bridges,  and  the  great  num- 
bers of  machines  made  from  iron  which  man  did  not  have  even 
a  century  ago.  Iron  ranks  first  among  the  metals  in  tonnage  used 
and,  next  to  aluminum,  it  is  the  most  abundant  metal  in  the  earth's 

Where  does  this  iron  come  from?  Iron  is  not  found  free,  except 
in  meteorites  that  have  fallen  on  the  earth  as  visitors  from  outer 
space.  Iron  ores,  however,  are  numerous  and  widely  distributed. 
The  most  important  iron  ore  deposits  (chiefly  hematite,  composed 
largely  of  Fe2O3)  occur  in  the  "ranges"  near  Lake  Superior.  The 
discovery  of  these  huge  deposits  of  iron  had  a  greater  effect  on  Ameri- 
can Iffe  than  the  more  romantic  discovery  of  gold  in  1849.  One  of 
the  iron  mines  in  the  Mesabi  range  yields  every  two  weeks  a  volume 
of  iron  equal  to  the  Great  Pyramid.  Here  the  ore,  containing  about 




50-55  percent  iron,  is  mined  in  open  cuts  by  power  shovels;  but  else- 
where underground  mining  is  generally  employed. 

These  Mesabi  range  mines  are  about  1000  miles  from  coke-  and 
steel-producing  centers.  In  the  metallurgy  of  iron  "the  ore  generally 
comes  to  meet  the  coal."  Why?  At  the  mines,  therefore,  the  ore  is 
loaded  into  railroad  cars  of  60-ton  capacity,  and  hauled  to  the  Great 
Lakes  port.  Here,  ore  boats  open  their  hatches,  and  a  load  of  10,000 
tons  of  ore  is  emptied  into  each  boat  in  half  an  hour.  After  the  long 
haul  down  the  Great  Lakes  to  lower  lake  ports,  huge  grab  buckets 
scoop  out  the  ore  and  drop  it  on  a  stock  pile  near  the  furnaces, 
which  transform  these  mountains  of  ore  into  rails  and  other  products. 

A  smaller  steel  industry  is  located  around  Birmingham,  Alabama, 
where  a  geological  revolution  once  laid  down  coal,  limestone,  and 
iron  ore.  This  region  supplies  about  ten  percent  of  our  steel. 

Next  in  importance  to  the  American  iron  ore  deposits  are  those  of 
the  Ruhr  and  of  Sweden,  which  also  help  to  supply  the  furnaces  of 
France,  Germany,  Belgium,  and  Luxembourg.  The  deposits  in  north- 
eastern England  and  Spain  are  valuable  also.  Russia,  too,  has  many 
rich  deposits  of  iron  ore.  Magnitogorsk,  a  city  in  the  Urals,  was 
built  in  1928  by  the  Soviet  Union  around  a  mountain  of  a  mag- 
netic iron  ore  called  magnetite,  Fe3O4.  Many  nations  do  not  have 
deposits  of  iron  ore.  They  depend  for  iron  on  other  nations. 

What  chemical  reactions  take  place  in  the  metallurgy  of  iron? 
The  metallurgy  of  iron  consists  essentially  of  the  reduction  of  iron 
oxide  by  means  of  carbon  monoxide  and  heated  coke.  This  reduc- 
tion takes  place  in  a  blast  furnace.  The  charge  that  enters  the  furnace 
consists  of  hematite,  coke,  and  flux.  The  hematite  supplies  the  iron; 
the  coke  (and  CO)  reduces  the  hematite;  and  the  flux  removes  the 
impurities  by  uniting  with  them,  forming  a  molten  slag  which  is 
drawn  off. 

of  Canada 

Surface  mining  of  iron  ore 
at  Seven  Islands,  Labrador. 
The  trucks  haul  the  ore  to 
a  nearby  railroad  line  to 
begin  the  long  trip  to  blast 
furnaces  in  the  United 




cold  blast 

(hot  gas) 

hot  gas  (part  to  heat  stoves) 

Adapted  from  drawing  by  American  Steel  and  Wire  Company 

Fig.  90.  Blast  furnace  and  stoves.  Study  this  diagram  in  conjunction  with  Fig.  91. 

The  nature  of  the  impurities  in  the  ore  determines  the  kind  of  flux 
used.  Oxygen  from  the  air  that  is  forced  through  the  charge  keeps  the 
coke  near  the  bottom  of  the  furnace  actively  burning.  Equations  for 
the  reactions  are  given  in  Fig.  91. 

A  blast  furnace  in  operation.  The  modern  blast  furnace,  which  has 
been  developed  from  a  furnace  invented  five  centuries  ago,  is  cylin- 
drical, about  30  feet  in  diameter  and  110  feet  high,  and  is  made  of 
steel  plates  lined  inside  with  firebrick,  a  refractory  substance  which 
withstands  very  high  temperatures.  At  the  top  of  each  blast  furnace 
is  an  apparatus  through  which  the  charge  may  be  dropped  with  a 
minimum  loss  of  heat.  About  eight  feet  from  the  base  are  12  pipes, 
or  tuyeres  (twfrs) ,  through  which  100  tons  of  hot  air  at  about 
600°C.  are  forced  every  hour.  The  heating  of  this  air  is  carried  out 
in  the  stoves,  cylindrical  towers  100  feet  high.  There  are  three  or 
four  stoves  to  each  blast  furnace. 

As  reduction  of  the  ore  proceeds,  the  lava-like  slag  floats  on  the 
molten  iron.  The  slag  is  removed  about  every  four  hours,  just  before 
the  iron  itself  is  tapped  and  poured  from  a  taphole  at  the  bottom  of 
the  furnace.  The  molten  iron  flows  out  along  a  gutter  lined  with 
firebrick  until  it  reaches  a  brick-lined  ladle  capable  of  holding  100 
tons  of  metal.  This  is  carried  to  an  endless  chain  conveyer  and 
poured  into  metal  forms.  The  early  forms  into  which  the  iron  flowed 
resembled  a  litter  of  feeding  pigs,  henc%  the/name  pig  iron. 

The  blast  furnace  workj  continuously  except  when  repairs  are 
necessary  or  business  is  slack.  Often  a  blast  furnace  is  in  operation 
day  and  night  for  many  months  at  a  stretch.  One  blast  furnace  pro- 
duces between  600  and  1200  tons  of  pig  iron  daily.  From  the  top  of 



the  furnace  more  than  3000  tons  of  hot  gases  issue  every  day.  These 
gases,  which  may  contain  as  much  as  25  percent  carbon  monoxide, 
are  gathered  and  used  to  heat  the  blasts  of  air  that  enter  the  tuyeres. 
Blast  furnace  gas  is  essentially  the  same  as  producer  gas. 

Composition,  properties,  and  uses  of  pig  iron.  Pig  iron  contains 
about  four  percent  carbon  and  smaller  amounts  of  compounds  of 
sulfur,  silicon,  and  phosphorus.  These  impurities  make  pig  iron  a 
grayish,  brittle  metal  which  melts  at  about  1150°C.  It  cannot  be 
forged  or  tempered,  but  can  be  cast  into  objects,  such  as  sash  weights, 
water  pipes,  stove  parts,  and  radiators,  that  do  not  have  to  withstand 
great  stresses  or  strains. 

Some  pig  iron  is  partially  purified  and  changed  to  cast  iron  by 
melting  it  with  scrap  or  wrought  iron  and  then  slowly  cooling  it.  Its 

Fig.  91.  Cross  section  of  a  blast  furnace.  Note  the  structure  of  the  furnace,  the 
nature  of  the  charge  and  the  chief  chemical  reactions  which  occur. 

Adapted  from  drawing  by  American  Iron  and  Steel  Institute 

waste  gases 

to  be  used 

to  heat  air  p^          I***|».J          P^l  ^  ore     f   char9e  to  be  dropped 

for  tuyeres 


lined  with 

hot  air 



taphole  for  pigiron 



carbon  is  present  as  globular  particles  of  graphite.  Such  cast  iron 
is  less  expensive  than  steel,  is  strong,  machinable,  and  can  be  welded. 
It  is  used  for  tools,  agricultural  instruments,  and  machine  parts  that 
are  not  subjected  to  severe  shocks. 

Most  pig  iron,  however,  is  converted  into  steel  by  removing  most 
of  its  impurities.  In  fact,  in  most  plants,  blast  furnaces  and  steel 
converters  stand  side  by  side. 

Wrought  iron  —  oldest  form  of  commercial  iron.  Up  to  the  four- 
teenth century  the  form  of  iron  most  commonly  used  was  wrought 
iron.  This  is  almost  pure  iron,  containing  less  than  0.1  percent  car- 
bon. Wrought  iron  is  soft  and  tough,  and  resists  shocks  and  strains 
very  well.  Although  it  has  been  almost  entirely  replaced  by  soft 
steel,  some  of  it  is  still  used  by  blacksmiths,  and  very  much  more  of 
it  is  used  in  the  manufacture  of  chains,  anchors,  pipes,  bolts,  and 
temporary  magnets.  It  resists  atmospheric  corrosion  rather  well. 

Wrought  iron  is  prepared  from  pig  iron  in  a  shallow  furnace  hold- 
ing about  600  pounds  of  the  melt.  Iron  oxide  or  rusty  scrap  iron  sup- 
plies the  oxygen  that  changes  the  carbon  of  the  pig  iron  to  carbon 

Fe2O3  4-  3C  ->  2Fe  -f  SCO  f 

As  the  iron  becomes  purer,  its  melting  point  rises  and  the  melt  be- 
comes a  pasty  mass.  This  is  stirred,  puddled,  and  then  removed  in  the 
form  of  fiery  balls  at  the  end  of  iron  rods.  These  balls  are  squeezed 
between  hammers,  which  remove  most  of  the  slag,  and  then  rolled 
to  distribute  the  remaining  slag  throughout  the  iron,  giving  the 
wrought  iron  a  fibrous  structure. 

Henry  Bessemer  and  William  Kelly  usher  in  the  steel  age.  Both 
wrought  iron  and  steel  have  been  known  for  centuries.  They  were 
made  in  small  amounts  by  skilled  workers  using  a  slow  and  expen- 
sive process  until  the  invention  of  the  rifled  cannon,  locomotive,  and 

Adapted  from  drawing  by  American  Iron  and  Steel  Institute 

*lfev;!/"     co' 


lining  of 
sand  bri 



Fig.    92.    Cross    section    of    a 
Bessemer  converter.  Note  the 
chemical    changes    which    oc- 
A       cur  as  the  air  passes  through 
slag  ^e  molten  iron. 

pig  iron 


other  machines  brought  an  unprecedented  and  immediate  demand 
for  huge  quantities  of  steel. 

William  Kelly,  a  Pittsburgher  living  in  Kentucky,  got  the  notion 
that  a  blast  of  air  would  not  chill  molten  pig  iron  but  would  increase 
its  heat  by  oxidizing  its  impurities.  In  this  way  he  planned  to  make 
steel  quickly  and  inexpensively.  He  was  scoffed  at,  and  his  father-in- 
law  actually  questioned  his  sanity  and  had  a  doctor  examine  him. 

Kelly  went  right  ahead,  however,  and  in  1851  he  gave  a  public 
demonstration  of  his  process.  An  eye-witness  reported:  "We  saw 
a  vessel  that  had  a  mouth  open  on  one  side  and  near  the  top.  The 
whole  was  shaped  something  like  an  egg.  We  saw  molten  metal 
poured  into  the  vessel.  Then  Kelly  turned  on  a  blast  of  cold  air.  The 
vessel  set  up  a  roaring  noise,  and  fire  belched  furiously  from  its 
mouth,  making  many  colors.  But  only  for  a  few  minutes.  The  noise 
and  fire  died  down.  We  then  saw  a  blacksmith  take  a  small  part  of 
the  iron  which  had  cooled,  and  contrived  and  threw  at  the  feet  of 
the  amazed  spectators  a  perfect  horseshoe.  No  one  laughed  at  Kelly 

While  Kelly  worked  to  improve  his  product,  Henry  Bessemer  in 
England  was  discovering  the  same  process  independently.  During 
the  Crimean  War,  he  had  invented  a  new  type  of  cannon  but  he 
could  not  find  an  iron  strong  enough  to  withstand  the  high  pressures 
of  the  expanding  gases  released.  This  led  him  to  researches  in  steel, 
and  in  1855,  at  the  age  of  42,  he  solved  the  problem  and  obtained 
a  patent  on  his  process.  Kelly  then  quickly  obtained  a  patent  for  his 
own  process  in  the  United  States.  Bessemer  built  a  more  efficient 
furnace  and  bought  out  Kelly's  patent. 

Bessemer  was  later  knighted.  This  honor  came  to  him  not  for  his 
invention  of  the  steel  furnace  but  rather  for  a  suggestion  he  had 
made  when  he  was  20  years  old.  He  had  recommended  an  improve- 
ment for  preventing  the  re-use  and  counterfeiting  of  seals  and  stamps 
on  official  documents. 

At  this  time  steel  was  selling  at  $250  a  ton.  The  world  was  wait- 
ing for  an  inexpensive  steel,  and,  almost  overnight,  the  iron  indus- 
try was  revolutionized.  The  first  Bessemer  furnace,  or  converter, 
used  in  the  United  States  was  set  up  at  Troy,  New  York,  in  1864. 
The  new  process  did  in  ten  minutes  what  the  old  process  took  a 
whole  month  to  accomplish.  The  old  puddling  furnaces  that  pro- 
duced wrought  iron  were  almost  completely  scrapped  and  soon 
steel  took  the  place  of  wrought  iron  almost  entirely. 

In  production,  pounds  were  replaced  by  tons.  By  1870  steel  pro- 
duction had  forged  ahead  so  rapidly  that  it  equalled  the  production 

One  of  the  most  spectacular  dis- 
plays in  industry— a  Bessemer 
converter  during  a  "blow." 

Kcthkhcm  Steel  Company 

of  wrought  iron.  At  present,  wrought  iron  represents  only  about  one 
percent  of  the  total  output  of  pure  iron  in  this  country.  The  highly 
skilled  ironworker  with  his  small  furnace  has  given  way  to  hundreds 
of  trained  men  handling  huge  furnaces  turning  out  steel  on  a  gigan- 
tic mass-production  scale. 

The  acid  Bessemer  process,  a  fiery  spectacle.  The  present-day  Bes- 
semer converter  used  in  the  United  States  is  a  barrel-shaped  pot  of 
wrought  steel  about  ten  feet  in  diameter  and  20  feet  deep,  lined  to 
a  thickness  of  1.5  feet  with  heat-resisting  brick  made  up  largely  of 
sand,  SiO,,  an  acid  anhydride.  Fifteen  tons  of  a  fiery  broth  of  mol- 
ten pig  iron,  whose  chief  impurities  are  carbon  and  silicon,  are 
poured  into  the  Bessemer  converter  from  the  blast  furnace. 

Through  this  molten  mass  blasts  of  cold  air  are  forced  under  about 
25  pounds  pressure,  through  holes  at  the  bottom  of  the  converter. 
This  blow  lasts  between  ten  and  20  minutes.  The  oxygen  in  the  blasts 
of  air  oxidizes  the  iron  it  strikes,  first  forming  ferrous  oxide,  which 
in  turn  reacts  with  the  silicon  impurity,  forming  silicon  dioxide  and 

2Fe  +  O2  -»  2FeO 
Si  +  2FeO  ->  SiO2  +  2Fe 

Silicon  dioxide,  insoluble  in  the  molten  iron,  accumulates  as  slag. 

As  the  blow  continues,  the  carbon  impurities  begin  to  burn.  A 
roaring  boil  then  takes  place  in  the  converter.  The  carbon  monoxide 



burns  at  the  mouth  of  the  converter,  and  countless  flying  sparks  of 
metal  and  slag  add  to  the  spectacle. 

C  +  O  >  —  CO2 
C02  +  C  -)  2CO 
2CO  +  02  -»  2CO2 

In  a  few  minutes  the  carbon  is  gone;  the  (lame  flickers  and  contracts. 
This  is  the  signal  to  stop  the  blast  by  turning  the  converter  vessel 
over  on  its  side. 

Inside  the  converter,  a  seething  mass  of  molten  iron  is  covered  with 
a  thin  layer  of  slag.  The  metal,  however,  contains  some  dissolved 
gases  and,  if  solidified  at  this  point,  would  be  spongy.  So  a  prede- 
termined amount  of  carbon  and  manganese  is  added  in  the  form  of 
iron  alloys,  such  as  ferromanganese,  which  contains  about  70  percent 
manganese,  or  spiegeleisen^  which  contains  about  20  percent  man- 
ganese. About  14  pounds  of  manganese  are  used  for  every  ton  of 
steel  produced.  The  manganese  unites  with  any  dissolved  air  or  com- 
bined oxygen,  and  in  addition  strengthens  the  steel  by  its  presence. 

The  chemical  reactions  that  take  place  in  the  converter  when 
spiegeleisen  is  added  are: 

<\  r™          ..         r^u  f  Mn  +  2FeO  — >  MnO2  +  2Fe 

1)  The  action  of  the  manganese:     |        Mn  +  ^  _  MnQ* 

2)  The  action  of  the  added  carbon: 

3Fe  -f  C  — >  Fe3C  (iron  carbide) 

The  converter  is  then  tipped  completely  over  and  its  charge  of  liquid 
steel  is  poured  into  a  waiting  mold. 

The  basic  open-hearth  process.  Pig  iron  containing  sulfur  and 
phosphorus  cannot  be  treated  in  the  acid  Bessemer  converter,  since 
this  type  of  furnace  cannot  remove  these  impurities.  Phosphorus  is 
oxidized  during  the  blow,  but  it  will  not  enter  the  slag  and  become 
part  of  it. 

The  basic  open-hearth  furnace,  however,  completely  removes  both 
sulfur  and  phosphorus  (acid  impurities)  by  means  of  lime  and  a 
lining  made  of  the  oxides  of  calcium  and  magnesium  which  are  basic 
anhydrides.  In  England  and  on  the  Continent  the  Bessemer  furnace 
lined  with  these  basic  oxides  is  used  in  the  treatment  of  iron  ores 
containing  sulfur  and  phosphorus. 

The  open-hearth  furnace  was  invented  by  Charles  W.  Siemens 
(ze'mens)  ,  who  left  Germany  to  become  a  citizen  of  England.  He 
patented  his  furnace  five  years  after  the  Bessemer  furnace  was  in- 
vented. In  1873  the  first  open-hearth  furnace  was  built  in  the  United 



Fig.  93.  Open-hearth  furnace 
(greatly  simplified)  showing 
the  nature  of  the  charge  and 
lining  and  the  flow  of  the  hot 

The  present  open-hearth  furnace,  as  large  as  an  eight-room  house, 
contains  a  shallow,  obiong  steel  basin  about  50  feet  long  by  15  feet 
wide  encased  in  brickwork.  The  charge  consists  of  about  equal 
amounts  of  scrap  iron  and  molten  pig  iron,  with  a  small  amount  of 
iron  ore  placed  over  a  layer  of  limestone,  CaCO:,.  The  scrap  iron  con- 
sists of  old,  discarded  factory  and  farm  machinery,  junked  cars,  and 
so  forth.  This  is  cheaper  than  pig  iron,  contains  less  carbon,  and 
conserves  our  reserves  of  raw  material.  Burning  fuel  and  hot  air 
enter  at  one  end,  are  deflected  down  from  the  low  roof  made  of  re- 
fractory brick,  and  heat  the  charge  just  as  a  gigantic  blast  lamp 
might;  the  waste  gases  finally  leave  at  the  other  end  of  the  furnace. 
Such  a  furnace  in  which  the  charge  is  heated  by  flames  deflected 
from  the  roof  is  called  a  reverberatory  furnace.  By  a  regenerative 
process,  all  the  heat  of  the  gases  is  used  by  reversing  their  flow  at 
regular  intervals.  The  following  reactions  take  place: 

1)  Carbon  is  burned  out  by  the  oxygen  of  the  air 

and  by  the  oxygen  of  the  rusty  scrap  iron,  and  ore. 
3C  +  Fe2O3  ~>  3CO  +  2Fe 

2)  The  limestone  decomposes,  forming  carbon  dioxide,  which 
bubbles  up  rapidly  through  the  melt.  This  bubbling  produces  a 
vigorous  stirring  and  leaves  a  basic  lining  of  CaO. 

CaC03  -»  CaO  +  CO2  1 

The  open-hearth  depart- 
ment of  a  steel  mill.  Open- 
hearth  furnaces,  located 
behind  the  wall  at  the  right, 
are  charged  with  molten 
iron  from  the  blast  furnace, 
scrap  iron,  and  limestone. 
About  12  hours  are  re- 
quired to  complete  a  "heat" 
of  250  tons  of  steel. 

A  special  machine  for  the  man- 
ufacture of  seamless  steel  tubing. 
Shown  is  a  length  of  white-hot 
tubing  just  after  piercing.  During 
the  manufacturing  process  a  steel 
billet  eight  feet  long  is  converted 
to  a  tube  35  feef  long. 

Rcpuhlic  Stcd  Corporation 

3)  Some  sulfur  is  liberated  as  sulfur  dioxide. 
' '-'••    <     v  -    :    ;-   "          FeS  +  2FeO  -*  3Fe  +  SO2 1 

4)  The  calcium  oxide  combines  with  the  phosphorus  and  silicon 
oxides,  forming  a  slag,  which  rises  to  the  top. 

3CaO  4-  P2O5  ->  Ca,(PO4)2 
CaO  H-  SiO2  ->  CaSiO3 

From  time  to  time,  samples  of  the  200-  to  500-ton  pool  of  liquid 
are  removed  and  quickly  analyzed  for  carbon  content  so  that  the 
composition  of  the  steel  may  be  controlled  carefully.  This  open- 
hearth  process  is  longer  and  may  be  controlled  more  easily  than  the 
acid  Bessemer  process,  and  takes  about  12  hours  for  200-ton  heats. 

In  spite  of  its  slightly  higher  price,  basic  open-hearth  steel  is  in 
great  dema